UNIVERSITY  OF  CALIFORNIA 
SAN  FRANCISCO  LIBRARY 


INORGANIC     CHEMISTRY 


WORKS  by  G.  S.  NEWTH,  F.I.C.,  F.C.S. 

DEMONSTRATOR   IN  THE   ROYAL  COLLEGE 
OF   SCIENCE,    LONDON. 


CHEMICAL  LECTURE  EXPERIMENTS. 
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CHEMICAL  ANALYSIS,  QUALITATIVE 
AND  QUANTITATIVE. 

With  100  Illustrations.     Crown  8vo,  $1.75. 

SMALLER  CHEMICAL  ANALYSIS. 
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A  TEXT-BOOK  OF  INORGANIC 
CHEMISTRY. 

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With  Appendix  of  Questions,  $2.00.    Appendix 

separately,  25  cents. 

ELEMENTARY  INORGANIC  CHEMISTRY. 

With  108  Illustrations  and  254  Experiments. 
Crown  8vo,  90  cents. 


LONGMANS,    GREEN,    AND    CO. 

NEW  YORK,  LONDON,  BOMBAY,  AND  CALCUTTA. 


A  TEXT-BOOK 


OF 


INORGANIC  CHEMISTRY 


G.  S.  NEVVTH,  F.I.C.,  RC.S. 
^^ 

SENIOR   DEMONSTRATOR-'IN   THE    ROYAL   COLLEGE   OF   SCIENCE    (IMPERIAL 

COLLEGE    OK    SCIENCE   AND   TECHNOLOGY1,    LONDON.       ASSISTANT- 

EXAMINER    IN   CHEMISTRY,    BOARD    OF    EDUCATION, 

SOUTH    KENSINGTON 


California  College  of  Pharmacy 


NEW   EDITION 


LONGMANS,    GREEN    AND    CO. 

91   AND  93  FIFTH   AVENUE,   NEW  YORK 
LONDON,  BOMBAY,  AND  CALCUTTA 


All  rights  reserved 


PREFACE 


IN  drawing  up  a  systematic  course  of  elementary  chemical 
instruction  based  upon  the  periodic  classification  of  the  ele- 
ments, whether  it  be  as  a  course  of  lectures,  or  as  a  text-book, 
a  number  of  serious  difficulties  are  at  once  encountered. 
These  possibly  are  sufficient  to  account  for  the  fact,  that 
although  twenty-five  years  have  elapsed  since  Mendelejeff 
published  this  natural  system  of  classification,  the  method  has 
not  been  generally  adopted  as  the  basis  of  English  elementary 
text-books. 

I  have  endeavoured  to  obviate  many  of  these  difficulties, 
while  still  making  the  periodic  system  the  foundation  upon 
which  this  little  book  is  based,  by  dividing  the  book  into 
three  parts.  Part  I.  contains  a  brief  sketch  of  the  funda- 
mental principles  and  theories  upon  which  the  science  of 
modern  chemistry  is  built.  Into  this  portion  of  the  book  I 
have  introduced,  necessarily  in  briefest  outlines,  some  of  the 
more  recent  developments  of  the  science  in  a  physico-chemical 
direction,  of  which  it  is  desirable  that  the  student  should  gain 
some  knowledge,  even  early  in  his  career. 

Part  II.  consists  of  the  study  of  the  four  typical  elements, 
hydrogen,  oxygen,  nitrogen,  and  carbon,  and  of  their  more 
important  compounds.  By  dissociating  these  four  elements 
from  their  position  in  the  periodic  system,  and  treating  them 
separately,  the  student  is  early  brought  into  contact  with  many 
of  the  simpler  and  more  familiar  portions  of  the  science.  Such 


vi  Preface 

subjects  as  water,  the  atmosphere,  and  combustion,  to  which  it 
is  desirable  that  he  should  be  introduced  at  an  early  stage  in 
his  studies,  are  thus  brought  much  more  forward  than  would 
otherwise  be  the  case. 

In  Part  III.  the  elements  are  treated  systematically,  accord- 
ing to  the  periodic  classification.  In  this  manner,  while 
avoiding  a  sharp  separation  of  the  elements  into  the  two  arbi- 
trary classes  of  metals  and  non-metals,  it  has  been  possible  to 
so  far  conform  to  the  prevailing  methods  of  instruction,  that 
all  those  elements  which  are  usually  regarded  as  non-metals 
(with  the  two  exceptions  of  boron  and  silicon)  are  treated  in 
the  earlier  portion  of  the  book. 

The  science  of  chemistry  has  of  recent  years  developed  and 
become  extended  to  such  a  degree,  that  the  difficulty  of  giving 
a  fairly  balanced  treatment  of  the  subject,  within  the  limits  of 
a  small  text-book,  is  an  ever-increasing  one,  and  it  necessarily 
resolves  itself  into  a  question  of  the  judicious  selection  of 
matter.  In  making  such  a  selection,  I  have  endeavoured,  as 
far  as  possible,  to  keep  in  view  the  requirements  of  students 
at  the  present  time,  without,  however,  following  any  examina- 
tion syllabus. 

Acting  upon  this  principle,  I  have  omitted  all  detailed 
description  of  the  rare  elements  and  their  compounds,  con- 
fining myself  merely  to  a  short  mention  of  them  in  a  few 
general  remarks  at  the  commencement  of  the  various  chapters. 

Although  from  a  purely  scientific  standpoint  many  of  these 
rare  substances  are  of  the  greatest  interest  and  importance, 
it  must  be  admitted  that  they  stand  quite  outside  the  range 
of  all  the  customary  courses  of  chemical  instruction ;  and  so 
far  as  the  wants  of  the  ordinary  student  are  concerned,  the 
space  which  would  be  occupied  by  an  account  of  these 
elements  is  more  advantageously  devoted  to  such  matters 


Preface  vii 

as  are  discussed  in  the  Introductory  Outlines.  Moreover,  it 
is  a  matter  of  common  observation  that  text-books,  even 
upon  the  shelves  of  reference  libraries,  and  which  bear  un- 
mistakable evidence  of  much  use,  are  frequently  uncut  in  those 
portions  which  treat  of  these  elements. 

Details  of  metallurgical  processes,  also,  are  out  of  place 
in  a  text-book  of  chemistry,  and  must  be  sought  in  metal- 
lurgical text-books.  Only  such  condensed  outlines,  therefore, 
have  been  given  as  are  sufficient  to  explain  the  chemicaj 
changes  that  are  involved  in  these  operations. 

The  great  importance  to  the  student,  of  himself  perform- 
ing experiments  illustrating  the  preparation  and  properties  of 
many  of  the  substances  treated  of  in  his  text-book,  cannot 
well  be  over-estimated.  If  he  be  in  attendance  upon  a  course 
of  chemical  lectures,  opportunity  should  be  given  to  him  for 
repeating  the  simpler  experiments  he  may  see  performed 
upon  the  lecture  table :  if  he  be  not  attending  lectures,  the 
necessity  for  this  practical  work  on  his  part  is  greater  still. 
Instead  of  burdening  this  text-book  with  specific  directions 
for  carrying  out  such  elementary  experiments,  frequent  refer- 
ences have  been  made  to  my  "Chemical  Lecture  Experi- 
ments," where  minute  directions  are  given  for  carrying  out 
a  large  number  of  experiments,  many  of  which  may  be  easily 
performed,  and  with  the  very  simplest  of  apparatus. 

Several  of  the  woodcuts  have  been  borrowed  from  existing 
modern  works,  such  as  Thorpe's  "  Dictionary  of  Applied 
Chemistry,"  Mendelejeffs  "Principles  of  Chemistry,"  Ost- 
wald's  "  Solutions,"  and  others.  Care  has  been  taken,  how- 
ever, to  exclude  all  antiquated  cuts,  and  a  large  number  of 
the  illustrations  are  from  original  drawings  and  photographs. 

G.  S.  N. 
SOUTH  KENSINGTON. 


HINTS  TO   STUDENTS 


FOR  the  help  of  students  who  may  use  this  book  at  the 
commencement  of  their  chemical  studies,  and  especially  for 
those  who  may  not  be  working  under  the  immediate  guidance 
of  a  teacher,  the  following  hints  are  given  : — 

Begin  by  carefully  reading  the  first  four  chapters  (pages 
1-24).  Then  pass  on  to  Part  II.  (page  171),  and  begin 
the  study  of  the  four  typical  elements,  hydrogen,  oxygen, 
nitrogen,  and  carbon,  and  their  compounds,  in  the  order  in 
which  they  are  treated.  Accompany  your  reading  by  per- 
forming as  many  of  the  experiments  referred  to  as  possible, 
in  order  that  you  may  become  practically  familiar  with  the 
substances  you  are  studying. 

During  the  time  occupied  in  the  study  of  these  four 
elements  and  their  compounds,  again  read  Chapters  I.  to  IV., 
and  slowly  and  carefully  continue  reading  Part  I.,  so  that 
by  the  time  Part  III.  is  reached,  you  may  have  fairly  mastered 
at  least  the  first  thirteen  chapters  of  the  Introductory  Out- 
lines. 

The  order  in  which  the  elements  are  treated  in  Part  III. 
is  based  upon  the  periodic  classification,  therefore  read  the 
short  introductory  remarks  at  the  commencement  of  the 
various  chapters,  in  the  light  of  the  table  on  page  118. 

Throughout  the  book  temperatures  are  given  in  degrees 
of  the  Centigrade  thermometer.  i°  Centigrade  equals  1.8° 


x  Hints  to  Students 

Fahrenheit,  and  as  the  zero  of  the  latter  scale  is  32°  below 
that  of  the  Centigrade,  temperatures  given  in  degrees  of  one 
scale  are  readily  translated  into  degrees  of  the  other,  by 
the  simple  formula  — 


The  abbreviation  mm.  stands  for  millimetre;  the  TTrVo-  Part 
of  a  metre  (i  1116^6  =  39.370113  inches;  or  roughly,  25 
mm.  =  i  inch).  The  abbrevation  c.c.  signifies  cubic  centi- 
metre; the  Y^IT  part  of  a  cubic  decimetre,  or  litre  (i 
litre  =  1.76077  pints). 

i  gramme  (the  weight  of  i  c.c.  of  distilled  water,  taken  at 
its  point  of  maximum  density)  =  15.43235  English  grains. 


TABLE    OF    CONTENTS 

PART  I 
INTRODUCTORY  OUTLINES 

CHAP.  PAGE 

I.  Chemical  Change — The  Constitution  of  Matter — Molecules — 

Atoms I 

II.  Elements   and   Compounds — Mixtures — Chemical   Affinity — 

Modes  of  Chemical  Action .6 

III.  Chemical  Nomenclature .15 

IV.  Chemical  Symbols         .         .         .         .         .         .         .         .21 

V.  The  Atomic  Theory — Laws  of  Chemical  Action     ...       25 

VI.  Atomic  Weights — Modes  of  Determining  Atomic  Weights      .  34 

VII.   Quantitative  Chemical  Notation    ......  53 

VIII.  Valency  of  the  Elements 59 

IX.  General  Properties  of  Gases — Relation  to  Heat  and  Pressure — 

Liquefaction — Diffusion — The  Kinetic  Theory         .         .  69 

X.   Dissociation — Reversible  or  Balanced  Actions        ...  88 

XI.  Electrolysis — Electrolytic  Dissociation — The  Ionic  Theory     .  96 

XII.   Classification  of  the  Elements — The  Periodic  System     .         .  112 

XIII.  General  Properties  of  Liquids — Evaporation,  Boiling,  Vapour 

Pressure  of  Solutions — The  Passage  of  Liquids  into  Solids 

— Freezing  Point  of  Solutions— Raoult's  Method     .         .126 

XIV.  Solution — Gases  in   Liquids — Liquids  in  Liquids — Solids  in 

Liquids — Osmotic  Pressure — Crystalline  Forms       .         .     142 
XV.  Thermo-chemistry 163 

PART   II 

THE   STUDY  OF  FOUR    TYPICAL   ELEMENTS 
Hydrogen  —  Oxygen  —  Nitrogen  —  Carbon, 

AND   THEIR   MORE    IMPORTANT   COMPOUNDS. 

I.  Hydrogen— Hydrogenium 171 

II.   Oxygen — Allotropy — Ozone 181 

III.  Compounds  of  Hydrogen  with  Oxygen 203 


xii  Contents 

CHAP.  PAGE 

IV.  Nitrogen 229 

V.  Oxides  and  Oxy-acids  of  Nitrogen 234 

VI.  The  Atmosphere  and  the  Argon  Group  of  Elements        .         .  252 
VII.  Compounds  of  Nitrogen   and    Hydrogen — Hydroxylamine — 

Ammon-sulphonates  ;  Halogen  Compounds  of  Nitrogen  .  272 

VIII.  Carbon 285 

IX.  Carbon  Monoxide — Carbon  Dioxide — Carbonates  .         .         .  295 
X.  Compounds  of  Carbon  with  Hydrogen — Methane — Ethylene— 

Acetylene 312 

XI.  Combustion — Heat  of  Combustion — Ignition  Point — Flame — 
Structure  of  Flame — Cause  of  Luminosity  of  Flames — 

The  Bunsen  Flame 321 


PART   III 

THE   SYSTEMATIC   STUDY   OF    THE    ELEMENTS,   BASED 
UPON   THE   PERIODIC   CLASSIFICATION 

I.  ELEMENTS  OF  GROUP  VII.  (FAMILY  B.) 

Fluorine:    Hydrofluoric   Acid.       Chlorine:    Hydrochloric 
Acid  —  Oxides   and    Oxyacids   of    Chlorine.      Bromine: 
Hydrobromic    Acid  —  Oxyacids    of    Bromine.      Iodine: 
Hydriodic  Acid—  Oxyacids  of  Iodine  —  Periodates  .         .     345 
II.  ELEMENTS  OF  GROUP  VI.  (FAMILY  B.) 

Sulphur:  Compounds  of  Sulphur  with  Hydrogen  —  Com- 
pounds with  Chlorine  —  Oxides  and  Oxyacids  of  Sulphur 
—  Oxychlorides  —  Carbon  Disulphide.  Selenium  —  Tellu- 


III. ELEMENTS  OF  GROUP  V.  (FAMILY  B.) 

Phosphorus:  Compounds  with  Hydrogen  —  Compounds  with 
the  Halogens  —  Oxides  and  Oxyacids.  Arsenic  :  Arsenu- 
retted  Hydrogen  —  Halogen  Compounds  —  Oxides  and 
Oxyacids  —  Sulphides.  Antimony:  Antimony  Hydride  — 
Halogen  Compounds  —  Oxides  and  Acids  —  Sulphides. 
Bismuth  :  Bismuth  and  Halogens  —  Oxides—  Sulphides  .  450 

IV.  ELEMENTS  OF  GROUP  I.  (FAMILY  A.) 

Potassium  —  Soditim  —  Lithium  —  Rubidium  —  Ammonium 
Salts  ..........  505 

V.  ELEMENTS  OF  GROUP  I.  (FAMILY  B.) 

Copper—  Silver—  Gold          .         .         ,         ,         ....         .     549 


Contents  xiii 

CHAP.  PAGE 

VI.  ELEMENTS  OF  GROUP  II.  (FAMILY  A.) 

Beryllium  —  Magnesiit  in  —  C ale  in  >n  — Strontium — Barium     5  70 

VII.  ELEMENTS  OF  GROUP  II.  (FAMILY  B.) 

Zinc — Cadmium — Mercury  590 

VIII.  ELEMENTS  OF  GROUP  III. 

FAMI LY  A. :    Scandium  —  ) 'tlrium  —  Lanthanum  —  Ytter- 
bium. 
FAMILY  B.  :    Boron — Aluminium — Gallium  —  Indium — 

Thallium 606 

IX.  ELEMENTS  OF  GROUP  IV. 

FAMILY  A.  :   Titanium — Zirconium — Cerium — Thorium. 
FAMILY  B.  :  Silicon — Germanium — Tin — Lead         .        .     627 

X.  ELEMENTS  OF  GROUP  V.  (FAMILY  A.) 

Vanadium — Niobium — Tantalum 655 

XL  ELEMENTS  OF  GROUP  VI.  (FAMILY  A.) 

Chromium — Molybdenum — Tungsten — Uranium         .         .     657 

XII.  ELEMENTS  OF  GROUP  VII.  (FAMILY  A.) 

Manganese 666 

XIII.  TRANSITIONAL  ELEMENTS  OF  THE  FIRST  LONG  PERIOD. 

Iron— Cobalt— Nickel 671 

XIV.  TRANSITIONAL  ELEMENTS  OF  THE  SECOND  AND  FOURTH 

LONG  PERIOD. 

Ruthenium  —  Rhodium — Palladiiim— Osmium — Iridium — 
Platinum — Argon — Helittm 690 

APPENDIX  :  RADIUM,  AND  RADIOACTIVE  ELEMENTS  .  697 

INDEX  7°5 


INORGANIC    CHEMISTRY 

PAET    I 
INTRODUCTORY    OUTLINES 


CHAPTER    I 
CONSTITUTION  OF  MATTER 

THE  science  of  chemistry  may  be  described  as  the  study  of  a 
certain  class  of  changes  which  matter  is  capable  of  undergoing. 
Matter  is  susceptible  of  a  variety  of  changes,  some  of  which  are 
regarded  as  physical  and  others  as  chemical.  Thus,  when  a  steel 
knitting-needle  is  rubbed  upon  a  magnet,  the  needle  undergoes  a 
change,  by  virtue  of  which  it  becomes  endowed  with  the  power 
of  attracting  to  itself  iron  filings  or  nails  ;  and  when  an  ordinary 
lucifer  match  is  rubbed  upon  a  match-box  the  match  undergoes  a 
change,  resulting  in  the  production  of  flame.  In  the  first  case  the 
change  is  said  to  be  a  physical  one,  while  the  ignition  and  com- 
bustion of  the  match  is  a  chemical  change. 

When  a  fragment  of  ice  is  gently  warmed,  it  is  changed  from  a 
hard,  brittle  solid  to  a  mobile,  transparent  liquid  ;  and  when  white 
of  egg  is  gently  heated,  it  changes  from  a  transparent,  colourless 
liquid  to  an  opaque  white  solid.  These  changes,  which  appear  at 
first  sight  to  be  of  a  similar  order,  are  in  reality  essentially  different 
in  their  nature  :  the  transformation  of  solid  ice  into  liquid  water 
is  a  physical  change,  the  coagulation  of  albumen  is  a  chemical 
change. 

Again,  when  certain  substances  (such  as  the  materials  which 
constitute  the  so-called  luminous  paint}  are  exposed  to  a  bright 
light,  they  undergo  a  change  whereby  they  become  invested  with 

A 


2  Introductory  Outlines 

the  power  to  emit  a  feeble  light  when  seen  in  the  dark.  A  stick  of 
phosphorus  also  emits  a  very  similar  light  when  seen  in  the  dark. 
The  glowing  of  these  materials  under  these  circumstances  might 
readily  be  regarded  as  the  result  of  the  same  kind  of  change  in 
both  cases  ;  but  in  reality  the  luminosity  of  the  phosphorus  is  due 
to  a  chemical  change  taking  place  upon  the  surface  of  that  sub- 
stance, while  the  emission  of  light  from  the  luminous  paint  is  a 
purely  physical  phenomenon. 

The  two  sciences,  chemistry  and  physics,  are  so  closely  related 
and  interdependent  upon  each  other,  that  no  sharp  distinction  or 
line  of  separation  between  them  is  possible.  Every  chemical 
change  that  takes  place  is  attended  by  some  physical  change,  and 
it  often  happens  that  this  accompanying  physical  change  forms 
the  only  indication  of  the  chemical  change  that  has  taken  place. 
In  certain  important  points,  however,  a  chemical  change  is  very 
different  from  one  that  is  purely  physical  :  in  the  latter  case  no 
material  alteration  in  the  essential  nature  of  the  substance  takes 
place.  This  will  be  seen  in  the  examples  quoted.  The  steel 
needle  remains  unaltered  in  its  essence,  although  by  magnetisation 
it  has  acquired  a  new  property — a  property  which  it  again  loses, 
and  which  can  be  again  and  again  imparted  to  it.  The  match,  on 
the  other  hand,  when  ignited  has  undergone  a  material  and  per- 
manent change  :  the  combustible  substance  is  now  no  longer 
combustible,  neither  will  it  ever  return  to  its  original  state.  The 
solid  water,  in  being  transformed  to  liquid  water,  has  not  under- 
gone any  vital  change  ;  in  essence  it  is  the  same  substance  merely 
endowed  with  a  new  property  of  liquidity,  a  property  which  it  loses 
again  when  cooled,  and  which  can  be  again  and  again  imparted  to 
it.  On  the  other  hand,  the  coagulated  albumen  has  undergone  a 
complete  and  lasting  change,  and  never  returns  to  its  original 
condition. 

In  the  same  way,  the  luminous  paint  gradually  ceases  to  emit 
light,  and  returns  to  its  original  state  ;  it  may  be  exposed  to  the 
influence  of  light,  when  it  once  more  acquires  the  property  of 
phosphorescence,  and  this  change  may  be  brought  about  indefi- 
nitely, without  altering  the  intrinsic  nature  of  the  substance.  The 
glowing  phosphorus,  on  the  other  hand,  is  gradually  changed  into 
a  white  substance,  which  escapes  from  it  as  a  smoke  or  fume  ;  in 
the  act  of  glowing  the  phosphorus  is  undergoing  a  process  of  slow 
burning,  and  if  allowed  to  remain  will  continue  glowing  and  burn- 
ing until  the  whole  of  it  has  disappeared  in  the  form  of  smoke. 


Molecules  3 

The  Constitution  of  Matter.  Molecules.— Matter  is  regarded 
by  the  chemist  and  physicist  as  being  composed  of  aggregations 
of  minute  particles  ;  every  substance,  whether  it  be  solid,  liquid, 
or  gaseous,  presents  the  appearance  to  his  mind  of  a  vast  number 
of  extremely  minute  particles.  To  these  particles  the  name  mole- 
cules ("  little  masses  ")  has  been  given.  The  particles  or  molecules 
of  any  particular  substance  are  all  alike  :  thus  in  sulphur  the 
molecules  are  all  of  one  kind,  while  in  water  they  are  all  of  another 
kind  ;  the  chemical  properties  associated  with  sulphur  are  the  pro- 
perties of  the  individual  sulphur  molecules,  while  those  belonging 
to  water  are  the  properties  of  the  molecules  of  that  substance.  All 
matter,  therefore,  is  to  be  conceived  as  having  what  may  be  called  a 
grained  structure.  The  actual  sizes  of  molecules  is  a  matter  which 
has  not  yet  been  determined  with  exactness  ;  they  are  orders  of 
magnitude  which  are  as  difficult  for  the  mind  to  grasp  on  account 
of  their  minuteness,  as  many  astronomical  measurements  are  by 
reason  of  their  vastness.  It  is  certain  that  their  size  is  less  than 
half  a  single  wave-length  of  light,*  and  that  therefore  they  are 
beyond  the  visual  limits  of  the  microscope.  Some  general  idea 
of  their  order  of  magnitude  may  be  gathered  from  Lord  Kelvin's 
calculation,  that  if  a  single  drop  of  water  were  magnified  to  the  size 
of  the  earth,  each  molecule  being  proportionately  enlarged,  the 
grained  appearance  which  the  mass  would  present  would  probably 
be  finer  than  that  of  a  heap  of  cricket-balls,  but  coarser  than  a 
heap  of  small  shot. 

It  will  be  evident,  therefore^  that  in  the  strictest  sense  matter  is 
not  homogeneous^  since  it  consists  of  aggregations  of  molecules, 
between  which  there  exist  certain  interspaces. 

The  forces  which  similar  molecules  exert  upon  each  other  are 
regarded  as  physical,  in  contradistinction  to  chemical.  These 
forces  are  either  attractive  in  their  nature,  or  repellent.  The 
attractive  forces  tend  to  draw  the  molecules  closer  together,  and 
thus  to  cause  the  substance  to  assume  the  solid  state  ;  while 
repellent  forces,  on  the  other  hand,  tend  to  separate  the  molecules 
and  to  make  the  substance  pass  into  the  gaseous  condition. 
Changes  which  matter  undergoes  by  the  action  of  these  forces  are 
physical  changes ;  they  do  not  affect  the  chemical  nature  and 
properties  of  the  substance,  which  properties,  as  already  stated, 
reside  in  the  molecules  themselves. 

*  The  wave-length  of  the  blue  ray  (G)  =  0.0004311  millimetre,  or 
0.0000169  inch. 


4  Introductory  Outlines 

In  each  of  the  three  states  of  matter,  viz.  solid,  liquid,  or  gaseous, 
the  molecules  are  conceived  as  being  in  a  state  of  motion  ;  they 
are  regarded  as  executing  some  vibratory  movement  within  the 
spaces  that  divide  them.  In  the  solid  state  this  movement  is 
usually  the  most  restricted,  for  the  reason  that  in  this  case  the 
intermolecular  spaces  are  as  a  rule  the  smallest.  In  the  gaseous 
condition,  however,  the  attractive  force  between  the  molecules 
has  been  almost  entirely  overcome  by  the  operation  of  the 
repellent  forces.  The  molecules  are  therefore  widely  sepa- 
rated, and  consequently  permit  of  a  much  greater  freedom  of 
movement. 

Such  changes  in  matter,  which  are  merely  the  result  of 
alterations  in  the  motions  of  the  molecules,  are  likewise  purely 
physical  changes. 

MoWi|]pq  may  HP  defined  as  the  smallest  weight  ofinatter 
in  which  the  original  properties  of  the  matter  are  retained. 

Atoms. — It  is  the  beliet  ot  chemists  that  most  molecules  are 
possessed  of  a  structure.  That  is  to  say,  they  are  not  simple, 
single,  indivisible  masses,  but  themselves  consist  of  aggregations 
of  still  smaller  particles,  which  are  held  together  by  the  opera- 
tions of  some  other  force.  These  particles  of  which  molecules 
are  composed  are  termed  atoms,  and  the  force  which  holds  them 
together  is  called  chemical  affinity,  or  chemical  attraction.  To 
the  mind  of  the  chemist,  such  molecules  are  little  systems,  con- 
sisting of  a  number  of  atoms  which  are  attracted  to  each  other 
by  this  particular  force  ;  in  the  ordinary  movements  of  the  mole- 
cule, the  system  moves  about  as  a  whole.  In  this  respect  it  bears 
some  analogy,  on  an  infinitely  minute  scale,  to  a  solar  system. 
The  atoms  of  a  molecule  are  regarded  as  in  a  state  of  motion  as 
respects  one  another,  possibly  revolving  about  one  another,  while 
the  entire  system,  or  molecule,  at  the  same  time  performs  its  in- 
dependent movements,  just  as  in  a  solar  system  the  various 
members  perform  various  movements  towards  each  other,  while 
at  the  same  time  the  whole  system  travels  upon  its  prescribed 
orbit.  In  the  case  of  the  heavenly  bodies,  the  force  which  regulates 
the  movements  of  the  individual  members  of  the  system  amongst 
themselves  is  the  same  force  that  controls  the  motion  of  the  united 
system,  namely,  gravitation.  What  is  the  precise  relation,  or 
difference,  if  any,  between  the  forces  which  control  the  move- 
ments of  molecules,  and  those  which  operate  between  the  atoms 
of  the  molecule,  is  not  known  ;  but  as  the  effects  produced  are 


Molecules  and  Atoms  5 

different,  the  latter  force  is  distinguished  by  the  name  of  chemical 
affinity. 

Any  change  which  matter  undergoes,  in  which  the  integrity  of 
the  molecules  is  not  destroyed,  is  regarded  as  a  physical  change  ; 
while  any  change  which  arises  from  an  alteration  in  the  structure 
of  the  molecule  is  a  chemical  change.  For  example,  the  molecules 
of  water  consist  of  three  separate  atoms,  one  of  oxygen  and  two 
of  hydrogen  ;  any  change  which  water  can  be  made  to  undergo, 
in  which  these  three  atoms  still  remain  associated  together  as  the 
molecule,  is  a  physical  change.  The  water  may  be  converted  into 
ice,  or  it  may  be  changed  into  steam  ;  but  these  alterations  still 
leave  the  molecules  intact  —  the  three  atoms  still  remain  united  as 
an  unbroken  system,  and  so  long  as  this  is  the  case  chemical 
change  has  not  taken  place. 

Suppose  now  the  molecules  of  water  are  heated  to  a  much 
higher  temperature  than  that  which  is  necessary  to  convert  the 
water  into  steam,  by  passing  electric  sparks  through  the  steam. 
It  will  then  be  found  that  a  very  different  kind  of  change  has  come 
over  the  substance.  The  steam,  after  being  so  heated,  no  longer 
condenses  to  water  again  when  cooled  ;  it  has  been  changed  into 
a  gas  which  can  be  bubbled  through  water  and  collected  in  an 
inverted  vessel  filled  with  water  standing  in  a  pneumatic  trough, 
and  if  a  flame  be  applied  to  this  gas  a  sharp  explosion  takes  place. 
The  change  in  this  case  is  a  chemical  change,  for  the  integrity 
of  the  molecules  of  water  has  been  destroyed.  The  two  atoms 
of  hydrogen  have  become  detached  from  the  oxygen  atom,  and 
the  original  triune  structure  of  the  system  is  destroyed. 

Atoms  are  therefore  defined  as  the  smallest  particles 
'which  can  take  part  in  a  chernical 


*  The  study  of  the  phenomena  of  radioactivity  has  led  to  the  belief  that 
atoms  are  not  indivisible  particles  of  matter,  but  that  they  are  themselves 
systems,  which  under  certain  circumstances  are  capable  of  undergoing  change 
by  ejecting  from  themselves  relatively  minute  portions  of  the  system  called 
electrons.  (See  p.  104;  also  Appendix.)  The  precise  nature  of  these  electrons 
still  belongs  to  the  realm  of  speculation,  and  the  changes  resulting  from  their 
movements  do  not  belong  to  the  category  of  "  chemical  change"  as  the  term 
is  here  employed. 


CHAPTER   II 
ELEMENTS  AND   COMPOUNDS 

THERE  are  certain  molecules  in  which  all  the  atoms  present  are 
of  the  same  kind,  and  there  are  other  molecules  which  are  com- 
posed of  atoms  which  differ  from  one  another.  Thus,  in  the 
substance  sulphur,  all  the  atoms  composing  the  molecules  are 
alike  ;  while  in  water,  as  already  mentioned,  there  are  two  distinct 
kinds  of  atoms  in  the  molecule.  Matter,  therefore,  is  divided  into 
two  classes,  according  as  to  whether  its  molecules  are  composed  of 
similar  or  of  dissimilar  atoms.  Molecules  consisting  of  atoms  of 
the  same  kind  are  termed  elementary  molecules,  and  substances 
whose  molecules  are  so  constituted  are  known  as  elements  ;  mole- 
cules, on  the  other  hand,  which  contain  dissimilar  elements  are 
called  compound  molecules,  and  substances  whose  molecules  are 
thus  composed  are  distinguished  as  compounds. 

Sulphur,  therefore,  is  an  element,  and  water  is  a  compound.  It 
will  be  evident  that  in  the  case  of  elementary  molecules,  whatever 
processes  they  may  be  subjected  to,  only  one  kind  of  matter  can 
be  obtained  from  them  ;  while  in  the  case  of  compounds,  the 
molecules  consisting  of  dissimilar  atoms,  as  many  different  kinds 
of  matter  can  be  obtained  as  there  are  different  atoms  present. 
By  appropriate  means  the  atoms  of  hydrogen  and  oxygen  in  water 
molecules  can  be  separated,  and  two  totally  different  kinds  of 
matter,  namely,  hydrogen  and  oxygen,  can  be  obtained  from  this 
compound. 

At  the  present  time  there  are  about  seventy  substances  known  to 
chemists  which  are  believed  to  be  elements.  In  the  history  of  the 
science  it  has  frequently  happened  that  substances  which  were 
considered  to  be  elements  have  proved,  when  subjected  to  new 
methods  of  investigation,  to  be  in  reality  compound  bodies  :  thus, 
prior  to  the  year  1783,  water  was  thought  to  be  an  elementary 
substance  ;  it  was  indeed  regarded  as  the  very  type  of  an  element, 
until  Cavendish  and  Lavoisier  proved  that  it  was  composed  of 
two  entirely  different  kinds  of  matter.  In  the  year  1807  Sir 


Elements  and  Compounds  7 

Humphry  Davy  showed  that  the  substances  known  as  potash 
and  soda,  which  were  believed  to  be  elements,  were  in  reality 
compound  bodies,  and  he  succeeded  in  separating  the  constituent 
atoms  in  the  molecules  of  these  substances,  and  in  obtaining  from 
them  two  essentially  different  kinds  of  matter.  It  is  therefore 
quite  possible,  perhaps  even  probable,  that  some  at  least  of  the 
forms  of  matter  which  are  now  held  to  be  elements  may  yet  prove 
to  be  compound  bodies.  On  the  other  hand,  the  list  is  from  time 
to  time  extended  by  the  discovery  of  new  elements.  Thus  during 
the  last  few  years  at  least  five  new  members  have  been  added  to 
the  number. 

The  number  of  compounds  is  practically  infinite. 

The  elements  are  very  unequally  distributed  in  nature,  and  are 
of  very  different  degrees  of  importance  to  mankind.  Some  are 
absolutely  essential  to  life  as  it  is  constituted,  while  others  might 
be  blotted  out  of  creation  without,  so  far  as  is  known,  their  absence 
being  appreciated.  The  following  thirty  elements  include  all  the 
most  important  (for  the  complete  list  see  page  22)  : — 

Aluminium.  Gold.  Oxygen. 

Antimony.  Hydrogen.  Phosphorus. 

Arsenic.  Iodine.  Platinum. 

Bismuth.  Iron.  Potassium. 

Bromine.  Lead.  Silicon. 

Calcium.  Magnesium.  Silver. 

Carbon.  Manganese.  Sodium. 

Chlorine.  Mercury.  Sulphur. 

Copper.  Nickel.  Tin. 

Fluorine.  Nitrogen.  Zinc. 

On  account  of  certain  properties  common  to  a  large  number  of 
the  elements,  and  more  or  less  absent  in  others,  properties  which 
are  for  the  most  part  physical  in  character,  the  elements  are 
divided  into  two  classes,  known  as  metals  and  non-metals.  The 
metals  generally  are  opaque,  and  their  smoothed  surfaces  reflect 
light  to  a  high  degree,  thus  giving  them  the  appearance  known  as 
metallic  lustre.  They  also  conduct  heat  and  electricity.  Gold, 
silver,  copper,  iron,  are  metals  ;  sulphur,  bromine,  oxygen,  phos- 
phorus, are  non-metals.  These  two  classes,  however,  gradually 
merge  into  one  another,  and  certain  elements  are  sometimes 
placed  in  one  division  and  sometimes  in  the  other,  depending 
upon  whether  the  distinction  is  based  more  upon  their  physical 


8  Introductory  Outlines 

or  their  chemical  properties  :  thus,  the  element  arsenic  possesses 
many  of  the  physical  properties  of  a  metal,  but  in  its  chemical 
relations  it  is  more  allied  to  the  non-metals  ;  such  elements  as 
these  are  often  distinguished  by  the  name  metalloids.  The  follow- 
ing list  embraces  all  those  elements  which  by  common  consent 
are  regarded  as  non-metals  and  metalloids,  including  the  recently 
discovered  elements  of  the  argon  group,  which  are  here  printed  in 
italics  : — 


Arsenic. 

Fluorine. 

Phosphorus. 

Helium. 

Boron. 

Hydrogen. 

Selenium. 

Neon. 

Bromine. 

Iodine. 

Silicon. 

Argon. 

Carbon. 

Nitrogen. 

Sulphur. 

Krypton. 

Chlorine. 

Oxygen. 

Tellurium. 

Xenon. 

The  number  of  atoms  which  compose  the  various  elementary 
•molecules  is  not  the  same  in  all  cases  :  thus  in  the  elements 
sodium,  potassium,  cadmium,  mercury,  and  zinc,  the  molecules, 
when  the  elements  are  in  a  state  of  vapour,  consist  of  only  one 
atom.  The  same  is  true  also  of  the  newly  discovered  elements  in 
the  last  column.  The  molecules  of  all  these  substances  are  single 
particles  of  matter.  The  terms  molecule  and  atom,  therefore,  as 
applied  to  these  elements,  are  synonymous.  Such  molecules  as 
these  are  called  mono-atomic  molecules.  In  many  cases  elemen- 
tary molecules  consist  of  two  atoms  ;  such  is  the  case  with  the 
elements  hydrogen,  bromine,  chlorine,  oxygen,  nitrogen,  and 
others.  Elementary  molecules  of  this  twin  or  dual  nature  are 
known  as  di-atomic  molecules.  Only  one  instance  is  known  in 
which  an  elementary  molecule  consists  of  a  trio  of  atoms,  namely, 
the  molecule  of  ozone,  which  is  an  aggregation  of  three  oxygen 
atoms.  This  molecule  is  said  to  be  tri-atomic.  In  two  cases, 
namely,  arsenic  and  phosphorus,  the  molecules  are  composed  of 
a  quartette  of  atoms,  and  these  elements,  therefore,  are  said  to 
form  tetr-atomic  molecules.  In  a  large  number  of  instances  the 
atomic  constitution  of  the  molecule  of  the  elements  is  not  known. 
These  terms,  mono-atomic,  di-atomic,  &c.,  are  applied  exclu- 
sively to  molecules  of  elements,  and  are  not  used  in  reference 
to  compounds,  where  the  molecules  are  composed  of  dissimilar 
atoms. 

Mechanical  Mixtures. — When  molecules  of  different  kinds  of 
matter  are  brought  together,  one  of  two  results  may  follow  :  either 
they  will  merely  mingle  together  without  losing  their  identity,  that 


Mechanical  Mixtures  9 

is  to  say,  the  atoms  composing  the  individual  molecules  will  still 
remain  associated  together  as  before,  or  the  atoms  in  the  molecules 
of  one  kind  will  attach  themselves  to  certain  atoms  present  in 
molecules  of  another  kind  to  form  still  different  molecules  ;  in  other 
words,  there  will  be  a  redistribution  of  the  atoms,  whereby  diffe- 
rent systems  or  molecules  are  produced. 

In  the  first  case  the  result  is  said  to  be  a  simple  or  mechani- 
cal mixture,  in  the  second  it  is  the  formation  of  a  chemical 
compound. 

In  a  simple  mixture  the  ingredients  can  be  again  separated  by 
purely  mechanical  methods  ;  and  as  the  properties  of  a  substance 
are  the  properties  of  the  molecules  of  that  substance,  it  follows  that 
if  the  integrity  of  the  molecules  is  not  broken,  the  properties  of  a 
mechanical  mixture  will  be  those  of  the  ingredients.  For  example, 
oxygen  is  a  colourless  gas  without  taste  or  smell ;  hydrogen  also  is 
a  colourless  gas  without  taste  or  smell :  when  these  two  gases  are 
mixed  together,  the  mixture  is  gaseous,  is  colourless,  and  tasteless, 
and,  being-  only  a  mixture,  the  molecules  of  one  gas  can  be  readily 
sifted  away  from  the  other. 

Again,  charcoal  is  a  black  solid,  insoluble  in  water  ;  sulphur  is  a 
yellow  solid,  also  insoluble  in  water  ;  nitre  is  a  white  solid,  readily 
dissolved  by  water :  when  these  three  substances  are  finely 
powdered  and  mixed  together,  the  result  is  a  mechanical  mixture, 
which  is  solid,  and  which  is  dark  grey  or  nearly  black  in  colour. 
If  this  mixture  be  placed  in  water,  the  nitre  is  dissolved  away  and 
the  charcoal  and  sulphur  are  left. 

When,  however,  the  integrity  of  the  molecules  is  disturbed,  and  a 
rearrangement  of  the  atoms  takes  place,  resulting  in  the  formation 
of  new  molecules,  then  it  is  said  that  chemical  action  has  taken 
place. 

Chemical  action,  therefore,  always  results  in  the  formation  of 
new  molecules — new  molecules  which  are  endowed  with  their 
own  special  properties,  differing  often  in  the  most  remarkable  and 
quite  inexplicable  manner  from  those  of  the  original  molecules. 
One  or  two  examples  may  be  quoted  in  order  to  illustrate  this 
extraordinary  modifying  effect  of  chemical  action.  The  two 
colourless  gases,  oxygen  and  hydrogen,  when  simply  mixed  to- 
gether, give  rise,  as  already  mentioned,  to  a  colourless,  gaseous 
mixture,  in  which  the  dual  molecules  of  hydrogen  and  the  simi- 
larly constituted  oxygen  molecules  move  about  freely  amongst 


IO  Introductory  Outlines 

each  other.  By  suitable  means  chemical  action  may  be  made 
to  take  place  between  these  two  elements,  whereby  a  complete 
rearrangement  of  the  atoms  takes  place,  resulting  in  the  formation 
of  molecules  of  water — molecules  in  which,  as  has  been  already 
mentioned,  one  atom  of  oxygen  is  associated  with  two  atoms  of 
hydrogen.  The  product  of  the  chemical  action  is  therefore  water, 
while  both  the  forms  of  matter  of  which  it  is  composed  are 
gaseous. 

The  air  we  breathe,  and  which  is  necessary  to  life,  consists  of 
a  simple  mixture  of  two  colourless  gases,  viz.,  oxygen  and  nitrogen. 
When  chemical  action  takes  place  between  these  substances,  a 
brown-coloured  gas  is  produced  in  which  no  animal  or  vegetable 
life  could  exist  for  many  minutes,  on  account  of  its  suffocating 
nature. 

Common  salt,  which  is  a  white  solid  substance,  and  not  only 
harmless,  but  even  a  necessary  article  of  food,  contains  two  atoms 
in  its  molecules — one  an  atom  of  chlorine,  and  the  other  an  atom 
of  sodium.  Chlorine  is  a  yellow  gas,  intensely  suffocating  and 
poisonous  ;  and  sodium  is  a  soft,  silver-like  metal,  which  takes 
fire  in  contact  with  water. 

Why  it  is  that  a  molecule,  consisting  of  an  atom  of  chlorine 
and  an  atom  of  sodium  held  together  by  chemical  affinity, 
should  be  endowed  with  properties  so  totally  different  from 
those  of  the  contained  elements,  is  altogether  unknown  ;  and 
similarly,  it  is  quite  impossible  to  predicate  from  the  properties 
of  any  compound  what  are  the  particular  elements  of  which  it  is 
composed.  Thus,  sugar  is  a  white  crystalline  solid,  soluble  in 
water,  and  possessing  a  sweet  taste  ;  but  no  one  would  have 
ventured  to  predict  that  the  molecules  of  this  substance  were  com- 
posed of  atoms  of  carbon,  a  black,  tasteless,  insoluble  solid  ; 
hydrogen,  a  colourless,  tasteless  gas ;  and  oxygen,  another  colour- 
less, tasteless  gas. 

Chemical  Affinity. — When  molecules,  consisting  of  two  atoms, 
say  A  B,  come  in  contact  with  molecules  consisting  of  other  two 
atoms,  C  D,  and  a  chemical  change  takes  place  resulting  in  the 
formation  of  new  molecules,  A  C  and  B  D,  the  question  naturally 
arises,  Why  does  the  atom  A  leave  the  atom  B  and  attach  itself  to 
C  ?  In  other  words,  what  determines  the  rearrangement  of  the 
atoms  into  new  molecules  ? 

At  present  no  exact  answer  can  be  given  to  this  question. 
Chemists  express  the  fact  by  saying  that  the  chemical  affinity 


Chemical  Affinity  1 1 

existing  between  A  and  C  is  greater  than  that  exerted  by  B  upon 
A.  This  remarkable  selective  power  possessed  by  the  atoms  of 
different  elements  lies  at  the  root  of  all  chemical  phenomena,  and 
it  differs  between  the  various  elements  to  an  extraordinary  degree. 
For  example,  the  atom  of  chlorine  possesses  a  very  powerful 
chemical  affinity  for  the  atom  of  hydrogen  :  when  hydrogen  mole- 
cules, which  consist  of  two  atoms,  are  mixed  with  chlorine  mole- 
cules, which  are  also  aggregations  of  two  atoms,  at  first  a  simple 
mechanical  mixture  is  obtained,  the  two  different  kinds  of  mole- 
cules move  amongst  each  other  without  undergoing  change.  On 
very  small  provocation,  however,  the  affinity  of  the  hydrogen  atoms 
for  the  chlorine  atoms  can  be  caused  to  exert  itself;  by  merely 
momentarily  exposing  the  mixture  to  sunlight  a  complete  redistri- 
bution of  the  atoms  suddenly  takes  place  with  explosive  violence 
and  new  molecules  are  formed,  each  containing  one  atom  of 
hydrogen  and  one  atom  of  chlorine. 

Again,  an  atom  of  nitrogen  is  capable  of  associating  itself  in 
chemical  union  with  three  atoms  of  the  element  chlorine,  forming 
a  compound  whose  molecules  therefore  contain  four  atoms.  The 
chemical  affinity  between  the  atoms  of  chlorine  and  nitrogen  is 
so  feeble,  the  system  is,  so  to  speak,  in  a  state  of  such  unstable 
equilibrium,  that  the  very  slightest  causes  are  sufficient  to  instantly 
separate  the  atoms  in  the  most  violently  explosive  manner,  and 
so  break  up  the  compound  molecules  into  separate  molecules  of 
chlorine  and  nitrogen.  In  this  case  the  affinity  between  one 
chlorine  atom  and  another  chlorine  atom  is  greater  than  that 
between  chlorine  and  nitrogen,  consequently  the  redistribution 
that  results  is  of  the  opposite  order  to  that  of  the  former 
example. 

As  a  rule,  those  elements  which  the  more  closely  resemble  each 
other  in  their  chemical  habits  have  the  least  affinity  for  each  other, 
while  the  greatest  affinity  usually  exists  between  those  which  are 
most  dissimilar. 

Chemical  Action. — The  actual  process  of  redistribution  of  the 
atoms  that  takes  place  when  molecules  of  different  kinds  of  matter 
are  brought  together  is  called  chemical  action.  In  many  cases 
chemical  action  takes  place  when  the  substances  are  merely 
brought  together,  while  in  others  it  is  necessary  to  expose  the 
bodies  to  the  influence  of  some  external  energy  :  thus  chemical 
action  is  brought  about  in  a  great  number  of  instances  by  the 
application  of  heat  to  the  substances.  In  some  cases  the  influence 


12  Introductory  Outlines 

of  light  has  the  effect  of  causing  chemical  action  to  take  place  ; 
for  example,  when  the  gases  chlorine  and  hydrogen  are  mingled 
together,  no  chemical  action  takes  place  in  the  dark,  but  on 
exposing  the  mixture  to  light  the  hydrogen  and  chlorine  combine, 
and  form  the  compound  hydrochloric  acid.  It  is  upon  the  effect 
of  light  in  causing  chemical  action  to  take  place  that  the  art 
of  photography  depends. 

Chemical  action  may  sometimes  be  induced  by  the  influence  of 
pressure  ;  thus,  when  the  two  gases,  hydrochloric  acid  and  phos- 
phoretted  hydrogen,  are  subjected  to  increased  pressure  they 
combine  together  to  form  a  crystalline  solid  compound  known  as 
phosphonium  chloride.  In  the  same  way,  by  very  great  mechanical 
pressure,  a  mixture  of  powdered  lead  and  sulphur  can  be  caused 
to  combine  together,  when  they  form  the  compound,  lead  sulphide. 
There  are  also  a  number  of  chemical  actions  that  are  only  able 
to  proceed  in  the  presence  of  small  quantities  (often  extremely 
small)  of  a  third  substance,  which  itself  remains  unchanged  at  the 
conclusion  of  the  action.  These  cases  are  generally  included 
under  the  name  of  catalytic  actions  :  in  some  of  them  the  modus 
operandi  of  the  third  substance  can  be  traced  (see  Oxygen,  Modes 
of  Formation  ;  also  Chlorine,  Deacon's  Process),  while  in  others 
it  is  not  understood.  Thus  it  is  found  that  a  number  of  chemical 
actions  are  quite  unable  to  take  place  if  the  materials  are  abso- 
lutely dry  ;  for  example,  the  element  chlorine  has  a  powerful 
affinity  for  the  metal  sodium,  and  when  these  substances  are 
brought  together  under  ordinary  conditions,  chemical  action  in- 
stantly takes  place,  and  the  compound  known  as  sodium  chloride 
(common  salt)  is  produced.  If,  however,  every  trace  of  moisture 
be  perfectly  removed  from  both  the  sodium  and  the  chlorine,  no 
action  between  these  elements  takes  place  when  they  are  brought 
together,  and  so  long  as  they  remain  in  this  state  of  perfect  dryness 
no  chemical  change  takes  place.  The  admission  into  the  mixture 
of  the  minutest  trace  of  the  vapour  of  water,  however,  at  once 
induces  chemical  action  between  the  chlorine  and  the  sodium,  but 
the  exact  part  that  the  trace  of  moisture  plays  in  producing  this 
effect  is  not  known  with  certainty.  (See  also  foot-note,  page  89.) 

A  few  interesting  cases  are  known  in  which  chemical  action  is 
brought  about  by  the  vibration  caused  by  a  loud  sound  or  note  ; 
for  example,  the  molecules  of  the  gas  acetylene  consist  of  two 
atoms  of  carbon  associated  with  two  of  hydrogen.  When  a  quantity 
of  this  gas  is  exposed  to  the  report  produced  by  the  detonation  of 


Chemical  Action  13 

mercury  fulminate,  the  mere  shock  of  the  explosion  causes  a  re- 
distribution of  the  atoms  whereby  solid  carbon  is  deposited  and 
hydrogen  set  free.  We  may  suppose  that  the  particular  vibration 
produced  by  the  detonation  of  the  fulminate  exercises  a  disturbing 
effect  upon  the  motions  of  the  atoms  constituting  the  molecules  of 
acetylene,  and  thereby  causes  them  to  swing  beyond  the  sphere  of 
their  mutual  attractions,  and  thus  the  system  undergoes  disruption 
and  rearrangement. 

All  known  instances  of  chemical  action  can  be  referred  to  one 
of  three  modes,  in  which  the  rearrangement  of  the  atoms  can  take 
place. 

(i.)  By  the  direct  union  of  two  molecules  to  form  a  more 
complex:  molecule.  Thus,  if  CO  and  C1C1  represent  two  mole- 
cules between  which  chemical  action  takes  place  according  to 
this  mode,  they  unite  to  form  a  molecule  containing  the  four 
atoms  COC1C1. 

(2.)  By  an  exchange  of  atoms  taking  place  between  different 
molecules.  In  its  simplest  form  this  is  illustrated  in  the  action 
of  one  element  upon  another  to  form  a  compound.  Thus,  if  HH 
and  C1C1  stand  for  two  elementary  molecules  between  which 
chemical  action  takes  place,  the  result  is  the  formation  of  the  two 
molecules  HC1  HC1.  Such  a  process  as  this,  in  which  a  com- 
pound substance  is  produced  directly  from  the  elements  which 
compose  it,  is  termed  synthesis. 

The  same  mode  of  chemical  action  may  also  be  exemplified  by 
the  exact  opposite  to  this  process,  namely,  the  resolution  of  a 
compound  into  its  constituent  elements.  Thus,  if  OHH  OHH 
represent  two  molecules  of  the  same  compound,  when  chemical 
action  takes  place  it  will  result  in  the  formation  of  the  three 
elementary  molecules  OO,  HH,  and  HH.  Such  a  process  as 
this,  in  which  a  compound  is  resolved  into  its  elements,  is  known 
as  analysis.* 

(3.)  By  a  rearrangement  of  the  atoms  contained  in  a  molecule. 
There  are  a  number  of  instances  of  chemical  change,  in  which  the 
molecules  of  the  substance  do  not  undergo  any  alteration  in  their 
composition — that  is  to  say,  no  atoms  leave  the  molecule,  nor  are 
any  added  to  it.  The  molecule  still  consists  of  the  same  atoms 
after  the  change  as  it  did  before,  but  the  chemical  action  has 

*  It  will  be  seen  that  in  each  of  the  examples  here  given,  the  process  of 
rearrangement  involves  first  the  decomposition  of  one  or  both  of  the  reacting 
molecules,  and  then  the  combination  of  the  atoms  to  form  different  molecules. 


14  Introductory  Outlines 

caused  them  to  assume  new  relative  positions,  or  different  relative 
motions  with  respect  to  each  other.  For  example,  the  substances 
known  to  chemists  as  ammonium  cyanate  and  urea  are  two  totally 
different  and  distinct  kinds  of  matter.  These  molecules,  however, 
each  contain  the  same  atoms  and  in  the  same  number  ;  they  each 
consist  of  aggregations  of  one  atom  of  carbon,  one  atom  of  oxygen, 
two  atoms  of  nitrogen,  and  four  atoms  of  hydrogen.  When  am- 
monium cyanate  is  gently  warmed,  the  eight  atoms  composing 
the  molecules  undergo  this  process  of  rearrangement,  and  the 
substance  is  changed  into  urea. 

When  chemical  action  takes  place  between  two  substances,  say  A  and  B, 
in  ordinary  language  we  say  that  A  acts  upon  B.  Such  a  statement,  however, 
must  not  be  understood  to  imply  that  A  takes  the  initiative,  so  to  speak,  and 
that  B  is  in  any  way  less  responsible  for  the  action.  It  is  equally  true  to  say 
that  B  acts  upon  A.  For  instance,  we  commonly  say  nitric  acid  acts  upon 
copper,  hydrochloric  acid  acts  upon  zinc,  nitric  acid  has  no  action  upon  gold, 
and  so  on ;  but  it  is  equally  true  to  say  copper  acts  upon  nitric  acid,  zinc 
acts  upon  hydrochloric  acid,  gold  has  no  action  upon  nitric  acid.  A  more 
strictly  scientific  expression  would  be  A  and  B  react,  or  do  not  react,  as  the 
case  may  be.  Thus,  nitric  acid  and  copper  react,  gold  and  nitric  acid  do  not 
react. 


CHAPTER   III 
CHEMICAL  NOMENCLATURE 

THE  names  which  have  been  given  to  the  various  elementary  forms 
of  matter  are  not  based  upon  any  scientific  system.  The  names  of 
some  have  their  origin  in  mythology.  Others  have  received  names 
which  are  indicative  of  some  characteristic  property,  while  those  of 
several  bear  reference  to  some  special  circumstance  connected  with 
their  discovery.  It  has  been  the  custom  in  modern  times  to  dis- 
tinguish metals  from  non-metals  by  applying  to  the  former  names 
ending  in  the  letters  um,  and  consequently  such  metals  as  are  of 
more  recent  discovery  all  have  names  with  this  termination.  The 
common  metals,  however,  which  have  been  known  since  earlier 
times,  such  as  gold,  silver,  tin,  and  copper,  keep  their  old  names. 
The  two  elements  selenium  and  tellurium  were  at  the  time  of  their 
discovery  thought  to  be  metals,  and  they  consequently  received 
names  with  the  terminal  um  ;  these  substances  strongly  resemble 
metals  in  many  of  their  physical  properties,  but  in  their  chemical 
relations  they  are  so  closely  similar  to  the  non-metal  sulphur,  that 
they  are  by  general  consent  classed  among  the  non-metals  ;  they  are 
examples  of  those  elements  which  are  distinguished  as  metalloids. 
On  this  account  selenium  is  by  some  chemists  termed  selenion. 

In  naming  chemical  compounds,  the  chemist  endeavours  that 
the  names  employed  shall  not  only  serve  to  identify  the  sub- 
stances, but  shall  as  far  as  possible  indicate  their  composition. 
The  simplest  chemical  compounds  are  those  composed  of  only 
two  different  elements  ;  such  are  spoken  of  as  binary  compounds* 
and  their  names  are  made  up  of  the  names  of  the  two  elements 
composing  them,  thus — 

*  This  expression  is  now  sometimes  used  in  a  somewhat  modified  sense. 
Thus  in  the  language  of  the  ionic  theory  (p.  107)  the  term  binary  compound  is 
used  to  denote  a  substance  which  dissociates  into  two  ions,  quite  irrespective 
of  the  number  of  elements  it  may  contain.  It  is  to  be  regretted  that  under 
these  circumstances  a  new  word  was  not  coined  to  denote  the  newer  idea. 


1 6  Introductory  Outlines 

The  compound  formed  by  the  chemical  union  of— 

Hydrogen  with  sulphur  is  called  hydrogen  sulphide. 
Sodium        ,,     chlorine       ,,        sodium  chloride. 
Copper        ,,     oxygen        ,,        copper  oxide. 
Calcium       ,,     fluorine        ,,        calcium  fluoride. 
Potassium    ,,     iodine  ,,        potassium  iodide. 

It  continually  happens,  however,  that  the  same  two  elements 
combine  together  in  more  than  one  proportion,  giving  rise  to  as 
many  different  compounds,  in  which  case  it  becomes  necessary  to 
so  modify  the  names  that  each  of  the  compounds  may  be  dis- 
tinguished. This  is  accomplished  by  the  use  of  certain  terminal 
letters  or  of  certain  prefixes ;  for  example,  the  element  phos- 
phorus combines  with  chlorine  in  two  proportions,  forming  two 
different  compounds — in  one  the  molecules  contain  one  atom  of 
phosphorus  united  to  three  atoms  of  chlorine,  in  the  other  the 
molecules  consist  of  one  atom  of  phosphorus  associated  with  five 
of  chlorine.  These  two  compounds  may  be  distinguished  in  the 
following  ways  : — 

i  atom  of  phosphorus  with  3  atoms  of  chlorine  forms  phosphorous  chloride. 
i         , ,  , ,  , ,          5         , ,  , ,  , ,          phosphorzV  chloride. 

or — 

i  atom  of  phosphorus  with  3  atoms  of  chlorine  forms  phosphorus  /rzchloride. 
i        ,,          ,,          ,,         5        ,,         ,,  ,,         phosphorus  /£«tachloride. 

The  latter  method  of  distinction  is  the  more  general,  thus — 

i  atom  of  sulphur  with  2  atoms  of  oxygen  forms  sulphur  dioxide, 
i         ,,         ,,         ,,         3        ,,          ,,          ,,          sulphur  trioxide. 
i  atom  of  carbon  with  i  atom  of  oxygen  forms  carbon  monoxide, 
i         ,,         ,,         ,,         2  atoms        ,,          ,,          carbon  dioxide. 

Occasionally  the  prefixes  sub  &&&proto  are  employed  to  denote 
these  differences  of  composition,  but  their  use  is  more  limited,  and 
is  becoming  out  of  vogue.  When  more  than  two  compounds  are 
formed  by  the  union  of  the  same  two  elements,  the  additional 
prefixes  hypo,  under,  and^ter,  over,  are  sometimes  used. 

In  a  considerable  number  of  instances  the  systematic  names  of 
familiar  compounds  give  way  to  the  vulgar  or  common  names  by 
which  they  are  known,  thus — 

{Ammonia      .        .         .  Hydrogen  nitride      ^ 

Hydrochloric  acid          .  Hydrogen  chloride     I  Systematic 

Sulphuretted  hydrogen .  Hydrogen  sulphide     j     names. 

Water  ....  Hydrogen  monoxide  ' 


Chemical  Nomenclature  17 

Binary  compounds  consisting  of  elements  united  with  oxygen 
are  called  the  oxides  of  those  elements.  Certain  of  these  oxides 
are  capable  of  reacting  with  water,  giving  rise  to  substances  known 
as  acids;  such  oxides  are  distinguished  as  acid-forming  oxides,  or 
acidic  oxides.  They  are  also  sometimes  termed  anhydrides.  All 
the  non-metallic  elements,  except  hydrogen  and  the  members  of 
the  argon  group,  form  oxides  of  this  order,  and  the  acids  derived 
from  them  are  known  as  the  oxy-acids  or  hydroxy-acids. 

Certain  other  oxides  also  unite  with  water,  but  give  rise  to  com- 
pounds known  as  hydroxides.  When  such  oxides,  which  are  all 
derived  from  the  metallic  elements,  are  brought  into  contact  with 
acids,  chemical  action  takes  place,  and  a  compound  termed  a  salt  is 
formed,  together  with  water.  Such  oxides  are  distinguished  as  salt- 
forming  or  basic  oxides.  There  are  also  oxides  which  are  neither 
acidic  nor  basic.  The  names  of  oxy-acids  are  derived  from  the 
name  of  the  particular  oxide  from  which  they  are  formed,  thus — 

Carbon  dioxide  gives  carbonic  acid. 
Silicon  dioxide      „      silicic  acid. 

When  the  same  element  forms  two  acid-forming  oxides,  the 
terminals  ic  and  ous  are  applied  to  the  acids  to  denote  respectively 
the  one  with  the  greater  and  the  less  proportion  of  oxygen,  thus — 

Sulphur  /rzbxide  gives  sulphur/V  acid. 
Sulphur  tfVoxide  gives  sulphurous  acid. 
Nitrogen  /<??;/oxide  gives  nitmr  acid. 
Nitrogen  /r/oxide  gives  nitrous  acid. 

When  more  than  two  such  acids  are  known,  the  additional 
prefixes  hypo  or  per  are  made  use  of.  Thus  jztersulphuric  acid 
denotes  an  acid  containing  the  highest  quantity  of  oxygen,  while 
fiyflonitrous  acid  stands  for  an  acid  containing  less  oxygen  than  is 
present  in  nitrous  acid. 

There  is  a  class  of  binary  compounds  formed  by  the  combination 
of  a  large  number  of  the  elements  with  sulphur  ;  these  are  known  as 
sulphides.  Certain  of  these  sulphides  are  also  capable  of -forming 
acids  which  are  analogous  in  their  constitution  to  oxy-acids,  but  in 
which  the  oxygen  atoms  are  substituted  by  atoms  of  sulphur. 
These  acids  are  known  as  thio  acids  (sometimes  sulpho  acids), 
and  the  same  system  of  nomenclature  is  adopted  to  distinguish 
these  :  thus  we  have  thio-arseni0«.y  acid,  thio-arsenzV  acid,  denoting 

B 


1 8  Introductory  Outlines 

respectively  the  acid  with  the  smaller  and  the  larger  proportion  of 
sulphur. 

It  was  at  one  time  believed  that  all  acids  contained  oxygen,  that 
indeed  this  element  was  essential  to  an  acid.  The  name  oxygen 
indicates  this  belief,  the  word  signifying  "the  acid-producer." 
This  view  is  now  seen  to  have  been  incorrect,  for  many  acids  are 
known  in  which  oxygen  is  not  one  of  the  constituents.  Thus  the 
elements  fluorine,  chlorine,  bromine,  and  iodine,  which  constitute 
the  so-called  Halogen  group  of  elements,  each  combines  with 
hydrogen,  giving  rise  respectively  to.  hydrofluoric,  hydrochloric, 
hydrobromic,  and  hydriodic  acids. 

All  known  acids  contain  hydrogen  as  one  of  their  constituents. 

As  already  stated,  when  chemical  action  takes  place  between  an 
acid  and  a  base  *  a  salt  is  formed.  Oxy-acids  in  this  way  give  rise 
to  oxy-salts,  thio-acids  to  thio-salts,  and  halogen  acids  to  haloid 
salts. 

The  latter  salts  being  binary  compounds,  their  names  are  given 
according  to  the  system  already  explained,  such,  for  example,  as 
calcium  fluoride,  sodium  chloride,  potassium  bromide,  silver  iodide. 

In  the  case  of  the  oxy-salts  and  thio-salts,  the  names  are  made 
up  from  the  names  of  the  acid  and  of  the  metal  contained  in  the 
base,  with  the  addition  of  certain  distinctive  suffixes  :  thus  if  the 
acid  be  one  whose  name  carries  the  terminal  ous,  its  salts  will  be 

*  The  word  base  is  unfortunately  employed  by  different  chemists  in  different 
senses,  so  that  it  is  scarcely  possible  to  give  a  precise  definition  of  it.  Originally, 
no  doubt,  the  term  was  employed  simply  to  denote  the  idea  of  foundation,  and 
was  applied  to  the  metal  or  the  oxide  of  the  metal  entering  into  the  composition 
of  a  salt ;  which  being  the  more  tangible  constituent  was  thus  regarded  as  the 
more  important  one,  or  the  basis  of  the  salt.  At  the  present  day  the  word 
base  is  used  in  INORGANIC  chemistry  chiefly  to  denote  that  class  of  compounds 
described  on  page.  17  as  hydroxides,  while  the  oxides  from  which  these 
hydroxides  are  derived  are  spoken  of  as  basic  oxides.  Besides  this  class,  it 
includes  ammonia  and  a  few  other  compounds  which  like  ammonia  are  not 
derived  from  metallic  oxides.  The  ORGANIC  chemist,  on  the  other  hand,  regards 
ammonia  as  the  true  type  of  a  base  ;  and  all  organic  compounds  which  can  be 
regarded  as  "  derivatives  "  of  ammonia  are  called  bases.  Not  only  so,  but  the 
term  is  even  extended  so  as  to  include  similar  "  derivatives  "  of  the  phosphorus, 
arsenic  and  antimony  analogues  of  ammonia,  thus  giving  rise  to  the  expressions 
nitrogen  bases,  phosphorus  bases,  &c. 

Again,  in  the  language  of  the  modern  theory  of  ionic  dissociation,  a  base  is 
defined  as  a  compound  in  which  the  only  negative  ions  are  the  hydroxide  ions 
(page  107).  This  definition  includes  the  class  of  hydroxides  above  mentioned, 
but  does  not  include  ammonia  gas. 


Chemical  Nomenclature  19 

distinguished  by  the  suffix  ite,  while  the  names  of  the  salts  derived 
from  acids  whose  names  end  in  ic  are  terminated  by  the  letters 
ate. 

Nitrous  acid  and  potassium  oxide  give  potassium  nitrz'/<?. 
Sulphurous  add  „  ,.  „         sulphite. 

Nitric  acid  „  „  ,,         nitrate. 

SulphurzV  acid  „  ,,  „         sulphate. 

The  formation  of  a  salt  by  the  action  of  an  acid  upon  a  base  is 
due  to  the  redistribution  of  the  atoms  composing  the  molecules  of 
the  two  compounds,  in  such  a  manner  that  some  or  all  of  the 
hydrogen  atoms  in  the  acid  molecules  exchange  places  with  certain 
metallic  atoms  from  the  molecules  of  the  base.  Acids  which  con- 
tain only  one  atom  of  hydrogen  so  capable  of  becoming  exchanged 
for  a  metal  are  termed  mono-basic  acids  ;  those  with  two,  three,  or 
four  such  hydrogen  atoms  are  distinguished  respectively  as  di-basic, 
tri-basic,  and  tetra-basic  acids. 

If  the  whole  of  the  displaceable  hydrogen  in  an  acid  becomes 
replaced  by  the  base,  the  salt  formed  is  known  as  a  normal  salt. 
On  the  other  hand,  when  only  a  portion  of  the  hydrogen  atoms 
is  displaced  by  the  base,  the  salt  is  distinguished  as  an  acia 
salt  Thus  sulphuric  acid  contains  two  atoms  of  hydrogen  in  its 
molecule  (associated  with  one  of  sulphur  and  four  of  oxygen)  ;  if 
both  the  hydrogen  atoms  are  exchanged  for  potassium,  the  salt 
obtained  is  normal  potassium  sulphate,  and  when  only  one  is  so 
replaced  the  salt  is  known  as  acid  potassium  sulphate.  By  the 
term  acid  salt,  therefore,  must  be  understood  not  a  substance 
having  the  familiar  properties  of  an  acid,  such  as  a  sour  taste  and 
the  power  to  redden  litmus,  but  a  salt  in  which  one  or  more  of  the 
hydrogen  atoms  of  the  original  acid  are  still  left  in  the  molecule.* 
It  is  quite  true  that  some  of  the  salts  of  this  class  do  possess  acid 
qualities  and  will  redden  litmus,  but  this  is  due  to  what  may  be 
regarded  as  merely  the  accidental  circumstance  of  the  acidic 
portion  of  the  molecule  being  derived  from  a  strong  acid.  Many 
substances  belonging  to  the  class  of  acid  salts  are  perfectly  neutral 
in  their  behaviour  towards  litmus,  while,  on  the  other  hand,  some 
are  strongly  alkaline.  For  example,  acid  potassium  sulphate  is  acid 

*  Some  chemists  prefer  to  regard  the  acids  themselves  as  the  hydrogen  salts ; 
accordingly  they  apply  to  nitric  acid,  sulphuric  acid,  nitrous  acid,  sulphurous 
acid,  &c. ,  the  names  hydrogen  nitrate,  hydrogen  sulphate,  hydrogen  nitrite, 
hydrogen  sulphite,  &c. ,  respectively. 


2o  Introductory  Outlines 

to  test  paper,  add  calcium  carbonate  is  neutral,  while  acid  sodium 
carbonate  is  alkaline. 

A  third  class  of  salts  is  formed  by  the  association  of  one  or 
more  molecules  of  normal  salt,  with  one  or  more  additional  mole- 
cules of  the  base  :  these  are  known  as  basic  salts.  Thus,  carbonic 
acid  and  the  base  lead  hydroxide  form  such  a  salt  known  as  basic 
lead  carbonate. 


CHAPTER   IV 
CHEMICAL    SYMBOLS 

CHEMISTS  are  agreed  in  adopting  certain  symbols  to  denote  the 
atoms  of  the  various  elementary  forms  of  matter.  The  table  on 
page  22  contains  the  names  of  the  elements  at  present  recognised, 
and  in  the  second  column  are  given  the  symbols  which  are  em- 
ployed to  represent  their  atoms.  The  names  of  the  rare  elements 
are  printed  in  italics. 

In  a  number  of  instances  the  atomic  symbol  is  the  initial  letter 
of  the  ordinary  name  of  the  element  :  thus  Boron,  B  ;  Carbon,  C  ; 
Fluorine,  F  ;  Hydrogen,  H  ;  Oxygen,  O  ;  Sulphur,  S. 

When  more  than  one  element  has  the  same  initial,  either  the 
first  two  letters  of  the  name,  or  the  first  and  another  that  is  pro- 
minently heard  in  pronouncing  the  word  are  employed,  as  Bromine, 
Br ;  Cobalt,  Co  ;  Chlorine,  Cl  ;  Platinum,  Pt.  In  some  cases 
letters  taken  from  the  Latin  names  for  the  elements  are  used,  such 
as  Antimony  (Stibium}^  Sb  ;  Gold  (Aurum\  Au  ;  Silver  (Argentuni), 
Ag  ;  Lead  (Plumbum\  Pb  ;  and  Iron  (Ferrum),  Fe. 

These  symbols  are  not  intended  to  be  employed  as  mere  short- 
hand signs,  to  be  substituted  as  abbreviations  for  the  full  names 
of  the  elements,  but  in  every  case  they  denote  one  atom  of  the 
element.  The  symbol  H  stands  for  one  atom  of  hydrogen,  the 
symbol  O  stands  for  one  atom  of  oxygen  ;  Cl  means  one  atom  of 
chlorine,  and  Ag  represents  one  atom  of  silver.  No  other  use  of 
these  symbols  is  legitimate. 

It  has  been  already  mentioned  (page  8)  that  the  molecules  of 
the  different  elements  are  composed  of  different  numbers  of  atoms  ; 
for  example,  the  molecule  of  hydrogen  consists  of  two  atoms,  and 
ordinary  oxygen  also  forms  diatomic  molecules.  These  facts  are 
expressed  in  chemical  notation  by  the  use  of  small  numerals  placed 
immediately  after  the  symbol  of  the  atom,  thus  H2  denotes  a  mole- 
cule of  hydrogen,  O2  a  molecule  of  oxygen.  The  molecule  of  ozone 
consists  of  an  aggregation  of  three  atoms  of  oxygen,  and  is 


22 


Introductory  Outlines 


Atomic  Weights. 
234 


Atomic  Weights. 
2          3 


1 

V 

rt 

_|1 

1 

<u 

ill 

E 

J  8 

•i|x 

g 

•i  !  . 

.2'cDp^ 

Name. 

OT 

X    3 
O  T3 

Name. 

C/3 

c 

'E 

P 

I'M 

o 

's 

!> 

\\i 

o 

<^ 

HH    B  "2 

o 

<^ 

^C     Q  VO 

4jj 

<  II 

«^ 

<  tt 

0 

0 

Aluminium 

Al 

27 

27.1 

Molybdenum   . 

Mo 

96 

96 

Antimony  (Sh'dinm) 

Sb 

120 

120.2 

Neodymium     . 

Nd 

144 

144-3 

Argon    .... 

A 

40 

39-9 

Neon      .... 

Ne 

20 

20.JL, 

Arsenic  .... 

As 

75 

75-° 

Nickel   .... 

Ni 

59 

58.7 

Barium  .... 

Ba 

137 

137-4 

Nitrogen     .     . 

N 

14 

I4.OI 

Beryllium  (Glucinum) 

Be 

9 

9.1 

Osmium      .     . 

Os 

191 

19*? 

Bismuth      .     . 

Bi 

208 

208.0 

Oxygen  .... 

O 

16 

16.00 

Boron     .... 

B 

11 

ii 

Palladium  .     .     . 

Pd 

106 

106.7 

Bromine 

Br 

80 

79.92 

Phosphorus 

P 

31 

31.0* 

Cadmium  .     .     . 

Cd 

112 

112.4 

Platinum    .     .     . 

Pt 

195 

1  95.X 

Caesium 
Calcium      .     . 

Cs 
Ca 

133 
40 

13** 

40.9* 

Potassium  (Kal-  ) 
ium]        .     .     ( 

K 

39 

39.10 

Carbon  .     . 

C 

12 

I2.OO 

Praseodymium    . 

Pr 

140-5 

I40.fr 

Cerium  .... 

Ce 

140 

140.25 

Radium       .     .     . 

Ra 

226-4 

226.4 

Chlorine     .     .     . 
Chromium  . 

Cl 
Cr 

35-5 
52 

35-4fc 
52.1 

i  Rhodium    . 
Rubidium  . 

Rh 
Rb 

103 
85 

103 
85-45 

Cobalt   .... 

Co 

59 

58-97 

Ruthenium 

Ru 

101-7 

101.7 

Columbium          ) 

(Niobium]    .     ) 

Cb 

93-5 

93-5 

Samarium 
Scandium  . 

Sm 
Sc 

150 
44 

150.4 
44.1 

Copper  (Cuprum] 
Erbium       .     . 

Cu 
Er 

63-5 
166 

63.6 

Selenium    .     .     . 
Silicon   .... 

Se 
Si 

79 
28 

79-2 
28.3 

Fluorine     .     .     . 

F 

19 

Z9~ 

Silver  (Argentum] 

Ag 

108 

107.88 

Gallium      .     , 

Ga 

70 

76 

Sodium  (Natrium 

Na 

23 

23.00 

Germanium     . 

Ge 

72 

72^5 

Strontium  . 

Sr 

87-6 

87.6 

Gold  (Aurum] 
Helium  .... 
Hydrogen  .     .     . 

Au 
He 
H 

197 
4 
1 

197(2 

4 
1.008 

Sulphur 
Tantalum  . 
Tellurium  . 

S 
Ta 
Te 

32 

181 
125  •? 

32-0^ 
127.6 

Indium  .... 

In 

115 

115 

Thallium    .     .     . 

Tl 

204 

204.0 

Iodine    .... 

I 

127 

126.92 

Thorium     .     .     . 

Th 

232 

232.4 

Iridium  .... 

Ir 

193 

I93-/ 

Tin  (Stannum]    . 

Sn 

118 

119 

Iron  (Ferrum] 

Fe 

56 

55-85 

Titanium  .     .     . 

Ti 

48 

48.1 

Krypton      .     .     . 

Kr 

81-5 

sjty. 

,  Tungsten    .     .     . 

W 

184 

184 

Lanthanum     . 

La 

139 

'39 

Uranium   . 

U 

238-5 

238.5 

Lead  (Plumbum] 

Pb 

207 

207.10 

Vanadium      .     . 

V 

51-2 

Lithium 

Li 

7 

7.0 

X 

128 

7^ft  <3^ 

Magnesium 

Mg 

24    . 

24.32 

Ytterbium  . 

Yb 

172 

172 

Manganese      .     . 

Mn 

55    X 

54-93 

Yttrium      .     .     . 

Y 

89 

89 

Mercury  (Hydr-  \ 
argyrum       .     j 

Hg 

200 

200.  fo 

Zinc  

Zirconium 

Zn 
Zr 

65 
90-7 

65.37 

90.6 

represented  by  the  symbol  O3,  while  the  tetr-atomic  character  of 
the  phosphorus  molecule  is  expressed  in  the  symbol  P4.  The 
composition  of  compound  molecules  is  expressed  by  placing  the 


Chemical  Symbols  23 

symbols  of  the  atoms  which  compose  such  molecules  in  juxta- 
position :  thus  a  molecule  consisting  of  one  atom  of  sodium  (symbol 
Na)  and  one  atom  of  chlorine  (symbol  Cl)  is  represented  by  the 
united  symbols  of  these  two  elements,  NaCl ;  a  compound  con- 
sisting of  one  atom  of  carbon  and  one  atom  of  oxygen  by  the 
symbols  of  these  two  atoms,  CO.  Such  arrangements  of  symbols 
representing  molecules  are  termed  molecular  formula,  or,  simply, 
formula. 

When  the  molecule  contains  more  than  one  atom  of  any  parti- 
cular element,  this  fact  is  indicated  by  the  use  of  numerals  placed 
immediately  after  the  symbol  to  be  multiplied  :  thus,  a  molecule  of 
water  consists  of  two  atoms  of  hydrogen  and  one  atom  of  oxygen; 
\\\Q  formula  for  water  is  therefore  H2O.  One  molecule  of  ammonia, 
consisting  of  an  atom  of  nitrogen  with  three  atoms  of  hydrogen, 
is  represented  by  the  formula  NH3  ;  and  a  molecule  of  sulphuric 
acid,  which  is  an  aggregation  of  two  atoms  of  hydrogen,  one 
atom  of  sulphur,  and  four  atoms  of  oxygen,  has  the  formula 
H2SO4. 

It  is  sometimes  necessary  to  represent  the  presence  in  a  mole- 
cule of  certain  groups  of  atoms,  groups  which  seem  to  hold  together 
and  often  to  function  as  a  single  atom.  This  is  accomplished  by 
the  use  of  brackets  :  thus  (NH4)2SO4  is  the  formula  for  a  molecule 
containing  one  atom  of  sulphur,  four  atoms  of  oxygen,  eight  atoms 
of  hydrogen,  and  two  atoms  of  nitrogen  ;  the  nitrogen  and  hydrogen 
atoms  being  present  as  two  groups,  in  each  of  which  one  nitrogen 
atom  is  associated  with  four  hydrogen  atoms.  Such  groups  of 
atoms  are  termed  compound  radicals. 

When  it  is  required  to  indicate  more  than  one  molecule  of  the 
same  substance,  numerals  are  placed  immediately  in  front  of  the 
formula  :  thus  2H2O  signifies  two  molecules  of  water,  and  3NH3 
expresses  three  molecules  of  ammonia. 

By  means  of  these  symbols  and  formulae,  chemists  are  enabled 
to  represent,  in  a  concise  manner,  the  various  chemical  changes 
which  it  is  the  province  of  chemistry  to  examine.  Such  changes 
are  usually  termed  chemical  reactions,  and  they  are  represented 
in  the  form  of  equations  in  which  the  symbols  and  formulae  of 
the  reacting  substances  as  they  are  before  the  change  are  placed 
on  the  left,  and  those  of  the  substances  which  result  from  the 
change  upon  the  right,  thus — 

H2  +  C12  =  2HC1 


24  Introductory  Outlines 

The  sign  +  has  a  different  significance  as  used  on  the  left  side 
of  the  equation  to  that  which  it  bears  upon  the  right.  On  the 
left  hand  it  implies  tfrat  chemical  action  takes  place  between  the 
substances,  while  on  the  opposite  side  it  has  the  simple  algebraic 
meaning.  Thus,  the  second  of  the  above  equations  is  understood 
to  mean,  that  when  the  compounds,  mercuric  chloride  and  potassium 
iodide,  are  brought  together  in  such  a  way  that  chemical  action 
results,  a  redistribution  of  the  atoms  will  take  place,  resulting  in 
the  formation  of  mercury  iodide  and  also  potassium  chloride. 

As  further  illustrations  of  the  use  of  chemical  symbols,  the 
following  three  examples  may  be  given  as  exemplifying  the  three 
modes  of  chemical  action  mentioned  on  page  13  : — 

(1)  NH3  +  HC1  =  NH4C1. 

Ammonia  combines  with  hydrochloric  acid,  and  gives  ammonium 
chloride. 

(2)  H2SO4  +  Na2CO3  =  Na2SO4  +  CO2  +  H2O. 

Sulphuric  acid  reacts  with  normal  sodium  carbonate,  and  yields 
normal  sodium  sulphate,  carbon  dioxide,  and  water. 

(3)  (CN)0(NH4)  =  (NH2)2CO. 
Ammonium  cyanate  is  converted  into  urea. 

In  all  cases  where  the  nature  of  the  chemical  change  is  under- 
stood, it  is  capable  of  expression  by  such  equations,  and  as  matter 
is  indestructible,  every  atom  present  in  the  interacting  molecules 
upon  the  left  of  the  expression  reappears  on  the  right-hand  side 
in  some  fresh  association  of  atoms.* 

*  See  also  Chemical  Notation,  chapter  vii. 


CHAPTER  V 

THE  ATOMIC  THEORY 

THE  atomic  view  as  to  the  constitution  of  matter,  briefly  sketched 
out  in  Chapter  I.,  forms  a  part  of  what  is  to-day  known  as  the 
atomic  theory. 

When  chemical  changes  were  carefully  studied  from  a  quantita- 
tive standpoint,  four  laws  were  discovered  in  obedience  to  which 
chemical  action  takes  place.  These  laws  are  distinguished  as 
the  laws  of  chemical  combination.  Three  of  these  generalisations 
refer  to  quantitative  relations  as  respects  weight ;  while  one  ex- 
presses quantitative  relations  with  regard  to  volume,  and  only 
relates  to  matter  in  the  gaseous  state. 

I.  Law  of  Constant  Proportion.— The  same  compound  always 
contains  the  same  elements  combined  together  in  the  same  proportion 
by  weight;  or  expressed  in  other  words,  The  weights  of  the  con- 
stituent elements  of  every  compound  bear  an  unalterable  ratio  to 
each  other,  and  to  the  weight  of  the  compound  formed. 

II.  Law   of  Multiple    Proportions.—  When   the   same   tiun. 
elements  confine  f^gtithflf  to  form,  more  thnn  nn^.   Compound,  tJie^ 
different  weights  of  one  of  the  elements  which  unite  with  a  constant 
weight  of  the  other  bear  a  simple  ratio  to  one  another:  or  this  law 
may  be  stated  thus :    \Vhen  one  element  unites  with  another  in 
two  or  more  different  proportions  by  weight,  these  proportions  are 
simple  multiples  of  a  common  factor. 

III.  Law  of  Reciprocal  Proportions,  or  Law  of  Equivalent 
Proportions. — The  weights  of  different  elements  which  combine 
separately  with  one  and  the  same  weight  of  another  element,  are 
either  the  same  as,  or  are  simple  multiples  of,  the  weights  of  these 
different  elements  which  combine  with  each  other;  or  in  other 
words,  The  relative  proportions  by  weight  in  which  the  elements, 
A,  B,   C,  D,  &*c.,  combine  with   a  constant  weight  of  another 
element,  X,  are  the  same  for  their  combinations  with  any  other 
element,  Y. 


26  Introductory  Outlines 

IV.  Law  of  Gaseous  Volumes,  or  The  Law  of  Gay-Lussae. 

—  When  chemical  action  takes  place  between  gases,  either  elements 
or  compounds,  the  volume  of  the  gaseous  product  bears  a  simple 
relation  to  the  volumes  of  the  reacting  gases. 

These  four  laws  are  the  foundations  upon  which  the  whole 
superstructure  of  modern  chemistry  rests. 

(i.)  The  Law  of  Constant  Proportions.— When  two  sub- 
stances are  mingled  together,  and  remain  as  a  mere  mechanical 
mixture,  they  may  obviously  be  present  in  any  proportion,  and  it 
was  at  one  time  thought  that  when  two  substances  entered  into 
chemical  combination  with  each  other,  they  could  do  so  also  in 
any  proportion,  and  that  the  composition  of  the  resulting  com- 
pound would  vary  from  this  cause.  This  belief  was  finally 
disproved,  and  the  law  of  constant  proportions  definitely  estab- 
lished by  Proust  in  the  year  1806.  The  same  compound,  therefore, 
however  made,  and  from  whatever  source  obtained,  is  always 
found  to  contain  the  same  elements  united  together  in  the  same 
proportion  by  weight.  Thus,  common  salt,  or,  to  adopt  its 
systematic  name,  sodium  chloride,  which  is  a  compound  of  the 
two  elements  sodium  and  chlorine,  may  be  made  by  bringing  the 
metal  sodium  into  contact  with  chlorine  gas,  when  the  two 
elements  unite  and  form  this  compound.  It  can  also  be  made 
by  the  action  of  hydrochloric  acid  upon  the  metal  sodium,  or  by 
adding  hydrochloric  acid  to  sodium  carbonate,  and  by  a  variety 
of  other  chemical  reactions.  When  the  sodium  chloride  obtained 
by  any  or  all  of  these  processes  is  analysed,  it  is  invariably  found 
to  contain  the  elements  chlorine  and  sodium  in  the  proportion  by 
weight  of  I  10.6479,  or,  expressed  centesimally — 

Sodium       .         .     39.32 
Chlorine     .         .     60.68 


100.00 

and  when  this  is  compared  with  the  sodium  chloride  as  found  in 
nature,  obtained  either  from  the  salt-mines  of  Cheshire,  or  the 
celebrated  mines  in  Galicia,  or  by  evaporating  sea-water,  it  is 
found  that  the  composition  of  the  compound  in  all  cases  is  exactly 
the  same.  In  the  same  way  the  compound  water,  consisting  of 
the  two  elements  hydrogen  and  oxygen,  whether  it  be  prepared 
synthetically  by  causing  the  two  elements  to  unite  directly,  or 
obtained  from  any  natural  source,  as  rain,  or  spring,  or  river,  is 


The  Atomic  Theory  27 

found  to  contain  its  constituent  elements  hydrogen  and  oxygen  in 
the  ratio  by  weight  of  i  :  8,  or, 

Hydrogen     .         .     11.12 
Oxygen         .         .     88.88 


IOO.OO 

If  in  the  formation  of  sodium  chloride  by  the  direct  combination 
of  its  constituent  elements,  an  excess  of  either  one  or  other  be 
present  beyond  the  proportions  39.32  per  cent,  of  sodium  and  60.68 
per  cent,  of  chlorine,  that  excess  will  simply  remain  unacted  upon. 
If  eight  parts  by  weight  of  hydrogen  and  eight  parts  by  weight 
of  oxygen  be  brought  together  under  conditions  that  will  cause 
chemical  action,  the  eight  parts  of  oxygen  will  unite  with  one  part 
of  hydrogen,  and  the  other  seven  parts  of  hydrogen  merely  remain 
unchanged.  This  fact,  that  elements  are  only  capable  of  uniting 
with  each  other  in  certain  definite  proportions,  marks  one  of  the 
most  characteristic  differences  between  chemical  affinity  and  those 
other  forces,  such  as  gravitation,  that  are  usually  distinguished  as 
physical  forces  ;  for  although  there  are  many  instances  known  in 
which  the  extent  to  which  a  chemical  action  may  proceed  (that  is, 
the  particular  proportion  of  the  reacting  bodies  which  will  undergo 
the  permutation  that  results  in  the  formation  of  different  mole- 
cules) is  influenced  by  the  mass  of  the  acting  substances,  it  never 
governs  the  proportion  in  which  the  elements  combine  in  these 
compounds. 

It  follows  from  the  law  of  constant  composition  that  the  sum  of 
the  weights  of  the  products  of  a  chemical  action  will  be  equal  to 
that  of  the  interacting  bodies ;  and  upon  the  validity  of  this  law 
depend  all  processes  of  quantitative  analyses. 

(2.)  The  Law  of  Multiple  Proportions  was  first  recognised 
by  Dalton,  who  investigated  certain  cases  where  the  same  two 
elements  combine  together  in  different  proportions,  giving  rise  to 
as  many  totally  distinct  compounds.  These  proportions,  however, 
were  always  found  to  be  constant  for  each  compound  so  produced, 
so  that  this  law  formed  no  contradiction  to  the  law  of  constant 
composition.  The  simple  numerical  relation  existing  between  the 
numbers  representing  the  composition  of  such  compounds  will  be 
evident  from  the  following  examples.  The  two'3*'  compounds  of 

*  In  Dalton's  day  these  two  substances  were  the  only  known  compounds  of 
carbon  with  hydrogen. 


28  Introductory  Outlines 

carbon   with  hydrogen,  known   as  marsh  gas  and  ethylene,   are 
found  to  contain  these  elements  in  the  proportions — 

Marsh  gas  .     .     i  part  by  weight  of  hydrogen  with  3  parts  of  carbon. 
Ethylene  i  ,,  »  »  6  ,,  ,, 

The  two  compounds  of  carbon  with  oxygen  contain  these  ele- 
ments in  the  proportion — 

Carbon  monoxide  .     i  part  of  carbon  with  1.334  parts  of  oxygen  by  weight. 
Carbon  dioxide  i  ,,  ,,  2.667  »>  »  » 

The  elements  nitrogen  and  oxygen  form  as  many  as  five  different 
compounds,  in  which  the  two  elements  are  present  in  the  propor- 
tions- 
Nitrous  oxide  .     .     i  part  of  nitrogen  with  o.  571  parts  of  oxygen  by  weight. 
Nitric  oxide.     .     .  ,,  ,,  1.143  ,.  »» 

Nitrogen  trioxide  .  ,,  ,,  i-7T4 

Nitrogen  peroxide  ,,  ,,  2.286  ,,  ,,  ,, 

Nitrogen  pentoxide  ,,  ,,  2.857 

The  relative  proportions  of  carbon  combining  with  a  constant 
weight  of  hydrogen  in  the  two  first  compounds  are  as  i  :  2. 

Those  of  oxygen  uniting  with  a  constant  weight  of  carbon  in  the 
second  example  are  also  as  i  :  2,  while  in  the  nitrogen  series  the 
relative  proportions  of  oxygen  in  combination  with  a  constant 
weight  of  nitrogen  are  as  i  :  2  :  3  :  4  :  5. 

(3.)  Law  of  Reciprocal  Proportions.— Known  also  as  the  law 
of  proportionality,  or  the  law  of  equivalent  proportions.  When 
the  weights  of  various  elements,  which  were  capable  of  uniting 
separately  with  a  given  mass  of  another  element,  were  compared 
together,  it  was  seen  that  these  weights  bore  a  simple  relation  to 
the  proportions  in  which  these  elements  combined  amongst  them- 
selves. For  example,  the  elements  chlorine  and  hydrogen  each 
.separately  combine  with  the  same  weight  of  phosphorus,  the  pro- 
portions being — 

Phosphorus  :  chlorine     =    i 
Phosphorus  :  hydrogen  =    i 

The  elements  chlorine  and  hydrogen  can  combine  together,  and 
they  do  so  in  the  proportion — 

Chlorine  :  hydrogen  =  35.5  :  i 

but  35  :  i  =  3-43  :  0.097 

Therefore   the   proportions  by  weight   in  which   chlorine  and 


The  Atomic  Theory  29 

hydrogen  separately  combine  with  phosphorus  is  a  measure  of  the 
proportion  in  which  they  will  unite  together. 

Again,  the  two  elements  carbon  and  sulphur  each  separately 
combine  with  the  same  weight  of  oxygen,  the  proportion  being — 

Oxygen  :  carbon    =  I  '.0.375 
Oxygen  :  sulphur  =  i  :  i 

But  the  elements  carbon  and  sulphur  themselves  unite  together, 
and  in  the  proportion — 

Carbon  :  sulphur  =  0.1875  :  r 
but    0.1875  :  :  =  °-375     :  2 

Therefore  the  proportion  by  weight  in  which  carbon  and  sulphur 
separately  unite  with  the  same  mass  of  oxygen  is  a  simple  multiple 
of  that  in  which  these  two  elements  combine  together.  These 
remarkable  numerical  relations  will  be  rendered  still  more  evident 
by  comparing  the  proportions  in  which  the  members  of  a  series  of 
elements  combine  with  a  constant  weight  of  various  other  elements  : 
thus — 

Hydrogen.     Sodium.     Potassium.     Silver.     Mercury.  Chlorine. 

0.02817      0.6479         1.02  3.04       2.816    unite  separately  with  i  part. 

It  will  be  seen  that  the  proportion  in  which  these  numbers  stand 
to  each  other  is  as — 

i       :       23       :       39      :      107      :      100  :  35.5 

Let  us  now  compare  these  proportions  with  those  in  which  the 
same  elements  unite  with  a  constant  weight  of  the  element 
bromine — 

Hydrogen.     Sodium.     Potassium.     Silver.     Mercury.  Bromine. 

0.0125        0.2875        0.4875        1.34         1.25      unite  with  i  part, 

or  as — 

i       :      23      :        39      :       107    :     100          :  80 

Each  of  these  five  elements  in  like  manner  combines  with 
oxygen,  and  the  weights  which  are  found  to  unite  with  a  constant 
mass  of  oxygen  are — 

Hydrogea     Sodium.     Potassium.     Silver.     Mercury.  Oxygen. 

0.125         2.875         4-875          I3-38        I2-5      unite  with  i  part, 

again  as — 

i       :        23       :       39      :       107    :     100  :  8 


30  Introductory  Outlines 

./" 

The  same  relation  will  appear  in  the  case  of  the  combination  of 
these  five  elements  with  a  constant  weight  of  sulphur — 

Hydrogen.     Sodium.     Potassium.     Silver.     Mercury.  Sulphur. 

0.0625        I-437S        2.4375        6.69        6.25      unite  with  i  part. 

or  as — 

i       :         23       :       39  107     :      100  :  16 

It  is  thus  evident  that  the  proportions  in  which  the  members  of 
such  a  series  combine  with  a  constant  weight  of  one  element  is  the 
same  as  that  in  which  they  unite  with  a  constant  mass  of  another 
element.  One  part  by  weight  of  hydrogen  combines  with  35.5 
parts  of  chlorine,  80  parts  of  bromine,  8  parts  of  oxygen,  and  16 
parts  of  sulphur — that  is  to  say,  these  proportions  of  these  four 
elements  satisfy  the  chemical  affinity  of  I  part  of  hydrogen  ;  they 
are  therefore  said  to  be  equivalent.  Twenty-three  parts  of  sodium 
is  likewise  equivalent  to  35.5  parts  of  chlorine,  80  parts  of  bromine, 
8  parts  of  oxygen,  and  16  parts  of  sulphur,  and  by  the  same 
reasoning  it  is  also  equivalent  to  i  part  of  hydrogen,  39  parts  of 
potassium,  107  parts  of  silver,  and  loo  parts  of  mercury.  These 
numbers,  therefore,  are  known  as  the  equivalent  'weights  or  the 
equivalents  of  the  elements,  or  their  combining  proportions,  and  the 
combining  weight  of  an  element  may  therefore  be  defined  as  the 
smallest  weight  of  that  element  which  will  combine  with  i  part  by 
weight  of  hvdrogen. 

This  law  of  proportionality,  or  reciprocal  proportion,  was  dis- 
covered by  Richter,  but  it  was  left  for  Dalton  to  trace  the  connec- 
tion between  these  three  generalisations.  Dalton  adopted  and 
adapted  an  ancient  theory  concerning  the  ultimate  constitution  of 
matter  which  was  expounded  by  certain  of  the  early  Greek  philo- 
sophers. The  exponents  of  this  theory  held  that  matter  is  built  up 
of  vast  numbers  of  minute  indivisible  particles,  in  opposition  to  the 
antagonistic  theory  believed  by  others,  namely,  that  matter  was 
absolutely  homogeneous  and  capable  of  infinite  subdivision. 

Dalton  embraced  the  ancient  doctrine  of  atoms,  and  extended  it 
into  the  scientific  theory  which  is  to-day  known  as  Dalton's  atomic 
theory,  and  is  accepted  as  a  fundamental  creed  by  modern  chemists. 

According  to  this  theory,  matter  consists  of  aggregations  of  minute 
particles,  or  atoms,  which  are  chemically  indivisible.  Dalton 
conceived  that  chemical  combination  takes  place  between  atoms — 
that  is  to  say,  when  chemical  action  takes  place  between  two 
elements,  it  is  due  to  the  union  of  their  atoms  ;  the  atoms,  coming 
into  juxtaposition  with  each  other  under  the  influence  of  chemical 


The  Atomic  Theory  31 

affinity,  are  held  together  by  the  operation  of  this  force.  He  further 
assumed  that  the  atoms  of  the  various  elements  possessed  different 
relative  weights,  and  that  the  relations  existing  between  these 
weights  was  the  same  as  that  between  the  weights  in  which  experi- 
ment had  shown  the  elements  to  be  capable  of  combining  together. 
In  other  words,  he  said  that  the  numbers  representing  the  combin- 
ing proportion  of  the  elements  expressed  also  the  relative  weights 
of  the  atoms. 

Let  us  now  see  how  this  theory  satisfies  and  explains  the  first 
three  laws  of  chemical  combination. 

(i.)  The  Law  of  Constant  Composition.— It  has  already  been 
shown  (p.  26)  that  the  compound  sodium  chloride,  wheresoever  and 
howsoever  obtained,  contains  the  elements  chlorine  and  sodium 
in  the  proportion — 

Chlorine  :  sodium  =  I  :  0.6479. 

These  numbers  have  been  shown  on  p.  29  to  represent  the  com- 
bining proportions — 

Chlorine  :  sodium  =  35.5  :  23. 

Now  the  atomic  theory  states,  that  sodium  chloride  is  formed  by 
the  union  of  atoms  of  chlorine  with  atoms  of  sodium,  and  that  the 
relative  weights  of  these  atoms  is  expressed  by  the  combining 
weights  of  the  elements,  namely,  35.5  and  23.  If  therefore,  sodium 
is  to  combine  with  chlorine,  since  atoms  are  indivisible  masses,  it 
follows  that  the  compound  produced  by  the  union  of  one  atom  of 
each  of  these  two  elements  must  always  have  the  same  composi- 
tion. 

(2.)  The  Law  of  Multiple  Proportions.— The  ratio  in  which 
oxygen  combines  with  hydrogen  to  form  the  compound  water  is 
seen  on  p.  27  to  be  as  8  :  i.  This  number  8,  therefore,  we  will 
for  the  present  argument  regard  as  the  relative  weight  of  the  atom 
of  oxygen.* 

Oxygen  combines  with  carbon  as  already  mentioned,  forming 
two  different  compounds  ;  in  the  first,  the  elements  are  present  in 
the  proportion — 

Carbon  :  oxygen  =  I  :  1.334  =  6:8, 

*  For  reasons  which  will  be  explained  later,  chemists  now  regard  the  number 
16  as  representing  (in  round  numbers)  the  relative  weight  of  the  atom  of 
oxygen. 


32  Introductory  Outlines 

that  is  to  say,  in  the  proportion  of  one  atom  of  carbon  to  one  atom 
of  oxygen.  According  to  the  theory,  if  the  atom  of  carbon  unites 
with  more  oxygen  than  one  atom,  it  must  at  least  be  with  two 
atoms.  It  may  be  with  three  or  with  four,  but  as  the  compound 
must  be  formed  by  the  accretion  of  these  indivisible  atoms,  the 
increment  of  oxygen  must  take  place  by  multiples  of  8.  When  the 
second  compound  is  examined  it  is  found  to  contain  its  constituent 
elements  in  the  proportion — • 

Carbon  :  oxygen  =i  :  2.667  =  6  :  16, 

that  is  to  say,  in  the  proportion  of  one  atom  of  carbon  to  two 
atoms  of  oxygen.  This  information  respecting  the  composition  of 
these  two  compounds  is  conveyed  both  in  their  names  and  their 
formulas.  The  first  is  termed  carbon  monoxide,  and  its  formula  is 
expressed  by  the  symbol  CO  ;  while  the  second  is  distinguished  as 
carbon  dioxide,  and  has  the  formula  CO2. 

The  difference  in  the  composition  of  the  five  compounds  that 
nitrogen  forms  by  union  with  oxygen  will  be  made  evident  by  the 
aid  of  this  theory.  The  proportion  of  nitrogen  to  oxygen  in  these 
compounds  is — 

(i.;  Nitrogen  :  oxygen  =  i  :  0.571  =  14  :  8 
(2.)  Nitrogen  :  oxygen  =  i  :  1.143  =  14  :  16 
(3.)  Nitrogen  :  oxygen  =  i  :  1.714  =  14  :  24 
(4.)  Nitrogen  :  oxygen  =  i  :  2.268  =  14  :  32 
(5.)  Nitrogen  :  oxygen  =  i  :  2.857  =  14  :  40 

And  it  will  be  seen  that  the  increase  in  the  proportion  of  oxygen  in 
the  compounds  takes  place  by  the  regular  addition  of  a  weight  of 
that  element  equal  to  8,  which  at  the  present  stage  of  the  argument 
we  are  regarding  as  representing  the  relative  weight  of  the  atom 
of  oxygen. 

(3.)  The  Law  of  Reciprocal  Proportions.— If  the  illustrations 
given  on  p.  28  of  the  operation  of  this  law  be  examined  in  the  light 
of  the  atomic  theory,  their  explanation  will  be  evident  :  thus,  the 
relative  proportions  in  which  hydrogen  and  chlorine  separately 
combine  with  phosphorus  is  0.097  :  3.43,  and  the  ratio  between  these 
numbers  is  as  I  :  35.5,  which  is  the  proportion  in  which  these  two 
elements  are  known  to  unite  together  to  form  hydrochloric  acid. 
These  numbers,  however,  represent  the  relative  weights  of  the 
atoms  of  these  elements,  therefore  hydrochloric  acid  may  be  sup- 
posed to  be  formed  by  the  union  of  one  atom  of  hydrogen  with 
one  atom  of  chlorine. 


The  Atomic  Theory  33 

Again,  the  relative  weights  of  carbon  and  sulphur  which  sepa- 
rately combine  with  a  constant  weight  of  oxygen  are — carbon,  0.375  ; 
sulphur,  i  ;  and  the  ratio  between  these  numbers  is  as  6  :  16. 

Carbon  and  sulphur,  however,  unite  together  in  the  relative 
proportion — 

Carbon  :  sulphur  =  0.1875  :  I  =  6  :  32. 

Therefore  the  compound  they  produce  may  be  supposed  to  consist 
of  one  atom  of  carbon,  having  the  relative  weight  6,  and  two  atoms 
of  sulphur,  each  with  the  relative  weight  16. 


CHAPTER   VI 
ATOMIC     WEIGHTS 

IN  the  third  column  of  the  table  on  page  22,  the  numbers  are 
given  which  are  at  the  present  time  generally  accepted  by  chemists 
as  representing  the  approximate  atomic  weights  of  the  elements. 
These  numbers  depart,  in  many  instances,  from  those  arrived  at 
by  Dalton's  methods  :  thus,  the  relative  weights  of  carbon,  oxygen, 
nitrogen,  and  sulphur,  which  were  found  to  be  equivalent  to  one 
part  of  hydrogen,  are — carbon  =  6,*  oxygen  =  8,  nitrogen  =  4.66, 
sulphur  =  1 6  ;  while  the  figures  given  as  the  approximate  atomic 
weights  of  these  elements  in  the  table  are — carbon  =  12,  oxygen 
=  1 6,  nitrogen  =  14,  sulphur  =  32.  We  must  now  discuss  some 
of  the  chief  reasons  for  these  departures.  In  the  two  compounds 
of  carbon  and  hydrogen  known  to  Dalton,  namely,  marsh  gas  and 
ethylene,  the  proportions  of  carbon  to  hydrogen  are — 

In  ethylene   .         .         .     Carbon  :  hydrogen  =  6  :  I. 
In  marsh  gas         .         .     Carbon  :  hydrogen  =  6:2. 

Dalton  therefore  concluded  that  ethylene  was  a  compound  con- 
taining i  atom  of  carbon  united  with  i  atom  of  hydrogen,  and  to 
which,  therefore,  he  gave  the  formula  CH  ;  and  that  marsh  gas 
consisted  of  I  atom  of  carbon  combined  with  2  atoms  of  hydrogen, 
and  which  he  accordingly  represented  by  the  formula  CH2. 

There  was,  however,  nothing  to  prove  that  the  weight  of  carbon 
was  constant  in  the  two  compounds,  for  it  will  be  obvious  that  the 
same  ratio  between  the  weight  of  carbon  and  hydrogen  will  still 
be  maintained  by  assuming  that  the  hydrogen  is  constant,  and 
that  the  carbon  varies,  thus — 

In  marsh  gas  .         .     Hydrogen  :  carbon  :   :  I  :  3. 
In  ethylene      .         .     Hydrogen  :  carbon  :   :  i  13x2. 

*  These  are  the  numbers  which  Dalton  ought  to  have  obtained  had  his 
methods  of  determination  been  more  exact.  The  figures  he  actually  found  for 
the  combining  weights  of  these  four  elements  were  respectively,  5,  7,  5,  13. 

34 


Atomic   Weights  35 

That  is  to  say,  the  ratios  are  not  disturbed  by  the  assumption 
that  in  marsh  gas  we  have  I  atom  of  hydrogen  combined  with  i 
atom  of  carbon,  having  the  relative  combining  weight  of  3,  and  in 
ethylene  I  atom  of  hydrogen  united  with  2  atoms  of  carbon. 

It  will  be  evident,  however,  that  if  we  could  gain  any  exact 
information  as  to  the  actual  number  of  atoms  which  are  present 
in  these  various  molecules,  this  difficulty  would  no  longer  exist. 

For  example,  suppose  it  were  possible  to  ascertain  that  in  the 
molecule  of  marsh  gas  there  were  4  atoms  of  hydrogen,  then  as 
the  relative  weights  of  hydrogen  and  carbon  in  this  compound  are 
as  I  :  3,  the  weight  of  the  carbon  atom  would  obviously  have  to 
be  raised  from  3  to  12  ;  and  if  it  could  be  determined  that  in  the 
ethylene  molecule  there  were  also  4  atoms  of  hydrogen,  then 
seeing  that  the  ratio  of  hydrogen  to  carbon  in  this  substance  is 
as  i  :  6,  we  should  conclude  that  it  contained  2  atoms  of  carbon, 
of  the  relative  weight  not  less  than  12,  and  the  composition  of  the 
two  compounds  would  be  expressed  by  the  formulae,  marsh  gas 
CH4,  ethylene  C2H4. 

Again,  the  relative  weights  of  hydrogen  and  oxygen  in  water 
are  as  i  :  8.  If  the  molecule  of  water  contains  only  i  atom  of 
hydrogen,  then  we  conclude  that  8  represents  the  relative  weight 
of  the  oxygen  atom,  and  the  formula  for  water  will  be  HO.  But 
suppose  it  to  be  discovered  that  there  are  two  atoms  of  hydrogen  in 
a  molecule  of  this  compound,  then  it  becomes  necessary,  in  order 
to  retain  the  ratio  between  the  weight  of  these  constituents  (a 
ratio  ascertained  by  analysis),  to  double  the  number  assigned  to 
the  oxygen  atom  and  to  regard  its  weight  as  16,  as  compared  with 
i  atom  of  hydrogen,  and  the  formula  for  water  in  this  case  would 
be  H2O. 

The  compound  ammonia  contains  the  elements  hydrogen  and 
nitrogen  in  the  ratio — 

Hydrogen  :  nitrogen  :  :  i  :  4.66. 

If  the  molecule  of  ammonia  contains  only  i  atom  of  hydrogen, 
then  4.66  represents  the  relative  weight  of  the  nitrogen  atom,  and 
the  formula  will  be  NH  ;  but  if  it  should  be  found  that  there  are 
3  atoms  of  hydrogen  in  this  molecule,  then  again  the  relative 
weight  assigned  to  the  nitrogen  must  be  trebled  in  order  to  pre- 
serve the  ratio,  and  it  will  have  to  be  raised  from  4.66  to  14  (in 
round  numbers),  and  the  formula  for  ammonia  will  be  NH3. 

From  these  considerations  it  will  be  evident,  that  it  is  of  the 


36  Introductory  Outlines 

highest  importance  to  gain  accurate  knowledge  as  to  the  actual 
number  of  atoms  which  are  contained  in  the  molecules  of  matter — 
in  other  words,  to  learn  the  true  atomic  composition  and  structure 
of  molecules  ;  and  it  may  be  said  that  this  problem  has  occupied 
the  minds  of  chemists  from  the  time  that  Dalton  published  his 
atomic  weights,  in  the  year  1808,  down  to  the  present  time.  There 
is  no  single  method  of  general  application,  by  means  of  which 
chemists  are  able  to  determine  the  atomic  weight  of  an  element  ; 
but  they  are  guided  by  a  number  of  independent  considerations, 
some  of  which  are  chemical  in  their  character,  while  others  are  of 
a  physical  nature  ;  and  that  particular  number  which  is  in  accord 
with  the  most  of  these  considerations,  or  with  what  are  judged  to 
be  the  most  important  of  them,  is  accepted  as  the  true  atomic 
weight. 

The  chief  methods  employed  for  determining  atomic  weights 
may  be  arranged  under  the  following  four  heads  : — 

1.  Purely  chemical  methods. 

2.  Methods  based  upon  volumetric  relations. 

3.  Methods  based  upon  the  specific  heats  of  the  elements. 

4.  Method  based  upon  the  isomorphism  of  compounds. 

I.  As  an  illustration  of  the  chemical  processes  from  which 
atomic  weights  may  be  deduced,  the  following  examples  may  be 
given,  namely,  the  case  of  the  two  elements  oxygen  and  carbon. 

Oxygen  combines,  as  already  stated,  with  hydrogen  in  the 
proportion — 

Hydrogen  :  oxygen  =  i  :  8. 

When  water  is  acted  upon  by  the  element  sodium,  the  compound 
is  decomposed  and  hydrogen  is  evolved ;  and  it  is  found  that  if 
1 8  grammes  of  water  are  so  acted  on,  I  gramme  of  hydrogen  is 
evolved,  and  40  grammes  of  a  compound  are  formed,  which 
contains  sodium,  together  with  all  the  oxygen  originally  in  the 
18  grammes  of  water,  and  some  hydrogen.  This  compound,  under 
suitable  conditions,  can  be  acted  upon  by  metallic  zinc,  and  when 
these  40  grammes  are  so  acted  on,  i  gramme  of  hydrogen  is  again 
evolved,  and  72.5  grammes  are  obtained  of  a  compound  containing 
no  hydrogen,  but  sodium  and  zinc  combined  with  all  the  oxygen 
originally  contained  in  the  18  grammes  of  water. 

It  will  be  evident,  therefore,  that  the  hydrogen  contained  in 
water  can  be  expelled  in  two  equal  moieties  ;  there  must,  therefore, 
be  two  atoms  of  hydrogen  in  this  compound.  By  no  known 


Atomic   Weights  37 

process  can  the  oxygen  be  withdrawn  from  water  in  two  stages  : 
thus,  if  1 8  grammes  of  water  are  acted  upon  by  chlorine,  under  the 
conditions  in  which  chemical  action  can  take  place,  73  grammes  of 
a  compound  containing  only  chlorine  and  hydrogen  are  formed,  and 
the  whole  of  the  oxygen  is  thrown  out  of  combination  and  evolved 
as  gas.  It  is  therefore  concluded  that  water  contains  in  its  mole- 
cule 2  atoms  of  hydrogen  and  I  atom  of  oxygen,  and  as  they  are 
combined  in  the  relative  proportion  of  I  :  8,  the  atomic  weight  of 
oxygen  cannot  be  less  than  16. 

No  compounds  have  been  found  in  which  a  smaller  weight  of 
oxygen,  relative  to  one  atom  of  hydrogen,  than  is  represented  by 
the  number  16  (approximately),  is  known  to  take  part  in  a  chemical 
change. 

The  compound  marsh  gas  contains  hydrogen  and  carbon  in 
the  proportion  by  weight  of  1:3.  By  acting  on  this  compound 
with  chlorine,  it  is  possible  to  remove  the  hydrogen  from  it  in 
four  separate  portions. 

By  the  first  action  of  chlorine  upon  16  grammes  of  marsh  gas, 
i  gramme  of  hydrogen  is  removed  in  combination  with  35.5 
grammes  of  chlorine,  and  a  compound  containing  carbon,  hydrogen, 
and  chlorine,  in  the  ratio  12  : 3  :  35.5,  is  formed. 

By  the  successive  action  of  chlorine,  three  other  moieties  of 
hydrogen  can  be  thus  withdrawn,  each  being  in  combination  with 
its  equivalent  (35.5  parts)  of  chlorine.  The  second  and  third  com- 
pounds that  are  formed  contain  carbon,  hydrogen,  and  chlorine  in 
the  ratios  12:2:  (35.5  x  2)  and  12:1:  (35.5  x  3). 

The  compound  produced  by  the  fourth  action  of  chlorine,  which 
withdraws  the  fourth  portion  of  hydrogen,  contains  only  carbon 
and  chlorine,  in  the  ratio  12  :(35-5  x  4).  From  the  fact  that  the 
hydrogen  contained  in  marsh  gas  can  thus  be  removed  in  four 
separate  portions,  the  molecule  must  contain  four  hydrogen  atoms, 
and  therefore  the  atomic  weight  of  carbon  must  be  at  least  12.  No 
compounds  of  carbon  are  known  in  which  a  smaller  weight  of 
carbon,  relative  to  one  atom  of  hydrogen,  than  is  represented  by 
the  number  12,  takes  part  in  a  chemical  change. 

The  definition  of  atomic  weight,  furnished  by  considerations 
of  a  chemical  nature,  may  be  thus  stated  :  the  atomic  weight  of  an 
element,  is  the  number  which  represents  how  many  times  heavier 
the  smallest  mass  of  that  element  capable  of  taking  part  in  a 
chemical  change  is,  than  the  smallest  weight  of  hydrogen  which 
can  so  function. 


38  Introductory  Outlines 

The  choice  of  hydrogen  as  the  unit  of  atomic  weights  is  a  purely  arbitrary 
selection  ;  but  since  atomic  weight  values  can  only  be  determined  relatively,  it 
becomes  necessary  to  select  some  one  element  and  to  assign  to  its  atom  some 
particular  number  to  serve  as  a  standard.  As  hydrogen  is  the  lightest  of  all 
elements,  Dalton  originally  adopted  it,  and  arbitrarily  fixed  unity  as  the 
number  which  should  stand  for  its  atomic  weight.  The  disadvantages  of  this 
particular  unit  are  twofold  :  in  the  first  place  the  number  of  elements  that  form 
hydrogen  compounds  that  are  suitable  for  atomic  weight  determinations  is  very 
small,  whereas  nearly  all  the  elements  form  convenient  oxygen  compounds,  or 
compounds  with  elements  whose  atomic  weights  with  reference  to  oxygen  are 
accurately  known,  and  in  actual  practice  such  compounds  are  almost  always 
made  use  of  for  such  determinations.  In  the  second  place,  the  exact  ratio  of 
the  weights  of  an  atom  of  hydrogen  and  oxygen  is  not  known  with  certainty,  so 
that  in  calculating  atomic  weights  that  are  determined  with  reference  to  oxygen, 
possible  errors  may  arise.  The  ratio  Hydrogen  :  Oxygen  is  not  exactly  i :  16. 
Various  values  have  been  obtained  by  different  experimenters,  and  at  the  present 
time  i :  15.88  is  accepted  as  more  nearly  the  truth. 

On  account  of  the  extreme  difficulty  of  exactly  determining  this  ratio, 
chemists  are  now  generally  agreed  in  adopting  as  the  unit  in  all  exact  determi- 
nations of  atomic  weights  a  number  which  is  -^th  the  weight  of  the  atom  of 
oxygen  :  that  is  to  say,  the  atomic  weight  of  oxygen  is  in  reality  the  standard, 
and  is  fixed  as  16,  and  the  unit,  instead  of  being  the  weight  of  one  atom  of 
hydrogen,  is  T^th  of  this  number. 

The  effect  of  this  change  is  only  of  importance  in  cases  of  chemical  investiga- 
tion where  a  high  degree  of  exactitude  is  required ;  for  purposes  of  ordinary 
analyses  and  chemical  calculations  the  difference  that  it  makes  is  practically  nil. 
Fixing  the  atomic  weight  of  oxygen  at  16  merely  raises  the  atomic  weight  of 
hydrogen  from  i  to  1.008.  As  the  use  of  small  decimal  fractions  introduces 
unnecessary  complications  which  tend  to  obscure  simple  processes  of  reasoning, 
the  approximate  atomic  weights  given  in  the  third  column  of  page  22  will  be 
employed  for  the  most  part  in  the  following  Introductory  chapters. 

The  student  will  frequently  meet  with  slight  discrepancies  between  the 
numbers  given  as  the  atomic  weights  of  various  elements  by  different  writers. 
Such  discrepancies  are  often  due  to  the  fact  that  in  some  cases  H  =  i  is  used 
as  the  standard,  and  in  others  O  =  16.  For  example,  the  atomic  weight  of 
gold  will  be  195.7  in  the  first  case,  and  197.2  in  the  second;  while  with  the 
lighter  metal  aluminium  the  numbers  will  be  26.9  as  against  27.1. 

The  discrepancy  may  also  arise  from  the  fact  that  the  determination  of 
atomic  weights  by  different  experimenters  often  vary  very  considerably.  With 
a  view  to  arrive  at  some  uniformity,  a  conference  of  representative  chemists 
was  held  to  consider  the  subject,  and  the  atomic  weights  finally  decided  upon 
by  them  were  published  under  the  title  of  International  Atomic  Weights.  A 
revised  list  of  these  weights  is  published  annually  in  the  Berichte,  and  in  the 
fourth  column  of  the  table  on  p.  22  will  be  found  the  latest  values  (1907). 

2.  Determination  of  Atomic  Weights  from  Considerations 
based  upon  Volumetric  Relations.  The  Law  of  Gaseous 
Volumes. — In  the  year  1805  the  fact  was  discovered  by  Gay- 
Lussac  and  Humboldt,  that  when  i  litre  of  oxygen  combines  with 


Atomic   Weights  39 

2  litres  of  hydrogen  the  vapour  of  water  (or  steam)  which  was 
produced  occupied  2  litres,  the  volumes  in  all  cases  being  measured 
under  the  same  conditions  of  temperature  and  pressure.*  This 
fact  led  to  the  discovery  of  the  simple  relation  existing  between 
the  volumes  of  other  reacting  gases  and  the  volume  of  the  products  : 
thus  it  was  found  that — 

I  vol.  of  hydrogen  unites  with   I    vol.  of  chlorine,  and  gives 
2  vols.  of  hydrochloric  acid. 

1  vol.  of  hydrogen  unites  with  I  vol.  of  bromine  vapour,  and 

gives  2  vols.  of  hydrobromic  acid. 

2  vols.   of  hydrogen  unite  with  I    vol.  of  oxygen,  and  give 

2  vols.  of  steam. 

2  vols.  of  carbon  monoxide  unite  with  I  vol.  of  oxygen,  and 
give  2  vols.  of  carbon  dioxide. 

1  vol.  of  carbon  monoxide  unites  with  i  vol.  of  chlorine,  and 

gives  i  vol.  of  phosgene  gas. 

In  the  same  way  with  compounds  that  cannot  be  obtained  by 
the  direct  union  of  their  constituent  elements,  it  is  found  that  on 
being  subjected  to  processes  of  decomposition  similar  simple 
volumetric  relations  exist :  thus  by  suitable  methods  of  decom- 
position— 

2  vols.  of  ammonia  gas  yield  i  vol.  of  nitrogen  and  3  vols.  of 

hydrogen. 
2  vols.  of  nitrous  oxide  yield  2  vols.  of  nitrogen  and  i  vol.  of 

oxygen. 
2  vols.  of  nitric  oxide  yield  i  vol.  of  nitrogen  and  I  vol.  of 

oxygen. 
I  vol.  of  marsh  gas  yields  2  vols.  of  hydrogen  and  some  solid 

carbon,  which   cannot   be   volatilised,  and   therefore   its 

vapour  volume  is  unknown. 
i  vol.  of  ethylene  yields  2  vols.  of  hydrogen  and  solid  carbon 

as  in  the  preceding. 

The  observations  of  these  and  similar  facts  gave  rise  to  the  law 
of  Gay-Lussac,  and  it  will  be  seen  that  there  is  evidently  a  close 
connection  between  the  simple  'volumetric  relations  and  those 
existing  between  the  multiple  proportions  by  weight,  in  which  one 

*  For  the  relations  of  gaseous  volumes  to  temperature  and  pressure  the 
student  is  referred  to  chapter  ix. ,  on  the  general  properties  of  gases. 


40  Introductory  Outlines 

element  unites  with  another.  For  example,  in  the  two  oxides  of 
nitrogen  the  ratios  of  the  two  elements  by  weight  are — 

Nitrous  oxide  .         .         .     Nitrogen  :  oxygen  =  28  :  16. 
Nitric  oxide     .         .         .     Nitrogen  :oxygen=i4  :  16, 

while  the  volumetric  relation  in  which  the  two  constituents  are 
present  is — 

Nitrous  oxide  .         .  .     Nitrogen  :  oxygen  =  2  :  I. 

Nitric  oxide      .         .  .     Nitrogen  :  oxygen  =i  :  i. 

In  other  words,  there  is  twice  as  much  nitrogen  by  weight  in  the 
one  compound  as  in  the  other,  and  there  is  twice  as  much  nitrogen 
by  volume  in  the  one  as  compared  to  the  other.  Moreover,  if  14 
and  1 6  respectively  represent  the  relative  weights  of  atoms  of  nitro- 
gen and  oxygen,  then  the  numbers  representing  the  relative 
volumes  in  which  these  elements  unite  will  also  express  the  number 
of  atoms  of  each  in  the  molecule. 

The  connection  existing  between  the  proportions  in  which 
elements  unite  by  weight,  and  by  volume,  was  first  explained  by 
the  Italian  physicist  and  chemist  Avogadro,  who  in  the  year 
1811  advanced  the  theory  now  recognised  as  a  fundamental  prin- 
ciple, and  known  as  Avogadro's  hypothesis.  This  theory  may  be 
thus  stated  :  Equal  volumes  of  all  gases  or  vapours,  under  the 
same  conditions  of  temperature  and  pressure,  contain  an  equal 
number  of  molecules.  If  this  be  true,  if  there  are  the  same 
number  of  molecules  in  equal  volumes  of  all  gases,  it  must  follow 
that  the  ratio  between  the  weights  of  equal  volumes  of  any  two 
gases  will  be  the  same  as  that  between  the  single  molecules  of  the 
particular  gases.  If  a  litre  of  oxygen  be  found  to  weigh  sixteen 
times  as  much  as  a  litre  of  hydrogen  (under  like  conditions  of  tem- 
perature and  pressure),  inasmuch  as  there  are  the  same  number 
of  molecules  in  each,  the  oxygen  molecule  must  be  sixteen  times 
heavier  than  that  of  hydrogen  ;  and  therefore  by  the  comparatively 
simple  method  of  weighing  equal  volumes  of  different  gases,  it 
becomes  possible  to  arrive  at  the  relative  weights  of  their  molecules. 

The  relative  weights  of  equal  volumes  of  gases  and  vapours,  in 
terms  of  a  given  unit,  are  known  as  their  densities  or  specific 
gravities.  Sometimes  densities  are  referred  to  air  as  the  unit,  but 
more  often  hydrogen,  as  being  the  lightest  gas,  is  taken  as  the 
standard.  Taking  hydrogen  as  the  unit,  the  density  or  specific 
gravity  of  a  gas  is  the  weight  of  a  given  volume  of  it,  as  compared 


Atomic   Weights  41 

with  the  weight  of  the  same  volume  of  hydrogen — or  in  other 
words,  the  ratio  between  the  weight  of  a  molecule  of  that  gas  and 
a  molecule  of  hydrogen.  The  ratio  that  exists  between  the  weight 
of  a  gaseous  molecule  and  half  the  weight  of  a  molecule  of 'hydrogen, 
chemists  term  the  molecular  weight  of  that  gas  ;  hence  it  will  be 
obvious  that  the  number  which  represents  the  molecular  weight  of 
a  gas  is  double  that  of  its  density  or  specific  gravity. 

If  i  litre  of  hydrogen  and  I  litre  of  chlorine  be  caused  to  combine, 
2  litres  of  gaseous  hydrochloric  acid  are  formed.  As  equal  volumes 
of  all  gases  (under  like  conditions)  contain  the  same  number  of 
molecules,  in  the  2  litres  of  hydrochloric  acid  there  must  be  twice 
as  many  molecules  of  that  compound  as  there  were  of  hydrogen 
molecules  in  the  I  litre,  or  of  chlorine  molecules  in  the  other. 
But  each  molecule  of  hydrochloric  acid  is  composed  of  chlorine 
and  hydrogen  (from  other  considerations  one  atom  of  each  element), 
therefore  there  must  have  been  at  least  twice  as  many  atoms 
of  hydrogen  in  the  litre  of  that  gas  as  there  were  molecules  ; 
and  by  the  same  reasoning,  twice  as  many  chlorine  atoms  in  the 
litre  of  chlorine  as  there  were  molecules  :  in  other  words,  both 
hydrogen  and  chlorine  molecules  consist  of  two  atoms.  The 
molecular  weight  of  hydrogen  therefore  is  2  ;  that  is,  its  molecule 
is  twice  as  heavy  as  its  atom.  The  atom  of  hydrogen  is  the  unit 
to  which  molecular  weights  are  referred,  while  the  weight  of  the 
molecule  of  hydrogen  is  taken  as  the  standard  of  densities  or 
specific  gravities. 

In  order,  therefore,  to  find  the  molecular  weight  of  any  gas  or 
vapour,  it  is  necessary  to  learn  its  density — that  is,  to  ascertain 
how  many  times  a  given  volume  of  it  is  heavier  than  the  same 
volume  of  hydrogen,*  and  to  double  the  number  so  obtained.t 

The  following  table  gives  the  densities  or  specific  gravities  of  all 
the  elements  whose  vapour  densities  have  been  determined.  The 
list  includes  all  those  elements  which  are  gases  at  the  ordinary 
temperature,  and  those  that  can  be  vaporised  under  conditions 

*  Certain  exceptions  to  this  rule  are  discussed  under  the  subject  of  Dissocia- 
tion, chap.  x.  p.  88. 

f  The  specific  gravity  of  hydrogen,  as  compared  with  air  taken  as  unity, 
is  0.0695,  or  air  is  14.3875  times  heavier  than  hydrogen.  If,  therefore,  it  be 
desired  to  find  the  molecular  weight  of  a  given  gas,  whose  density  as  compared 
with  air  is  known,  it  is  only  necessary  to  multiply  its  density  (air=i)  by  the 
number  14.3875,  which  gives  its  density  as  compared  with  hydrogen,  and  then 
to  double  the  number  so  obtained. 


42  Introductory  Outlines 

which  render  such  determinations  experimentally  possible.  (Hy- 
drogen being  taken  as  unity,  the  other  numbers  are  the  approxi- 
mate values,  which  for  purposes  of  discussion  are  more  suitable 
than  figures  that  run  to  two  or  three  decimal  places.) 


Hydrogen  I 

Helium  ...       2 

Neon  .  .  .  .10 

Nitrogen  .  .  .14 

Oxygen  .  .  .16 

Fluorine  .  .  19 

Argon  .  .  .20 

Sulphur  .  .  -32 

Chlorine  .  .  .     35.5 

Krypton  .  .  -41 


Selenium     ...  79 

Bromine     ...  80 

Iodine          .         .         .  127 

Sodium       .         .         .  11.5 

Potassium  .         .         .  19.5 

Zinc     ....  32.5 

Cadmium    ...  56 

Mercury      .         .         .  100 

Phosphorus          .         .  62 
Arsenic       .         .         .150 


Xenon         .         .         .     64.0 

Let  us  now  consider  how  the  "knowledge  of  the  relative  weights 
of  gaseous  molecules  is  utilised  in  assigning  a  particular  number 
as  the  atomic  weight  of  an  element. 

The  molecular  weight  of  chlorine  is  71.  It  has  been  shown  that 
the  molecule  certainly  contains  more  than  I  atom,  and  probably  2, 
in  which  case  35.5  would  represent  the  relative  weight  of  the 
atom. 

The  compound  hydrochloric  acid  has  the  molecular  weight  36.5. 
It  has  been  already  proved  that  this  compound  contains  I  atom  of 
hydrogen,  therefore  36.5  —  I  =35.5. 

The  compound  carbon  tetrachloride  gives  a  molecular  weight 
154.  Analysis  shows  that  this  compound  contains  12  parts  of 
carbon  in  154  parts,  therefore  154—12=142  =  35.5x4. 

In  these  three  molecules  the  weights  of  chlorine  relative  to  the 
weight  of  i  atom  of  hydrogen  are  142,  35.5,  and  71,  the  greatest 
common  divisor  of  which  is  35.5.  This  number,  therefore,  is 
selected  as  the  atomic  weight  of  chlorine. 

Again,  it  has  been  shown  that  by  the  action  of  metals  upon 
water,  the  hydrogen  contained  in  the  water  could  be  expelled  in  two 
separate  portions,  thus  proving  that  there  must  be  2  atoms  of 
hydrogen  in  the  molecule  of  that  compound. 

The  molecular  weight  of  water  is  found  to  be  1 8  ;  deducting  from 
this  the  weight  of  the  two  hydrogen  atoms  we  get  18  —  2  =  16. 

The  molecular  weight  of  carbon  monoxide  is  28  ;  28  parts  of 
this  compound  contain  12  parts  of  carbon,  therefore  28-12  =  16. 


Atomic  Weights  43 

The  molecular  weight  of  carbon  dioxide  is  44  ;  44  parts  of  this 
compound  also  contain  12  parts  of  carbon,  therefore  44-12  =  32. 

When  i  litre  of  oxygen  combines  with  two  litres  of  hydrogen, 
2  litres  of  water  vapour  are  formed  ;  there  are  therefore  twice  the 
number  of  water  molecules  produced  as  there  are  oxygen  mole- 
cules (since  by  Avogadro's  hypothesis  2  litres  contain  twice  as  many 
molecules  as  i  litre).  But  each  water  molecule  contains  certainly 
i  atom  of  oxygen,  therefore  the  original  oxygen  molecules  must 
have  consisted  of  not  less  than  2  atoms.  When  the  density  of 
oxygen  is  determined  it  is  found  to  be  16,  its  molecular  weight 
therefore  is  32. 

In  these  four  various  molecules  the  weights  of  oxygen  relative  to 
the  weight  of  I  atom  of  hydrogen  are  16,  16,  32,  32,  the  greatest 
common  divisor  of  which  is  16.  This  number,  therefore,  is  selected 
as  the  atomic  weight  of  oxygen. 

Again,  it  has  already  been  shown  that  in  the  compound  ammonia, 
the  hydrogen  can  be  removed  in  three  separate  moieties,  proving 
that  there  must  be  three  atoms  of  that  element  in  the  molecule. 
The  molecular  weight  of  ammonia  is  found  to  be  17,  therefore 
17  —  3  =  14,  which  is  the  weight  of  the  nitrogen. 

The  molecular  weight  of  nitrous  oxide  is  44  ;  44  parts  of  this 
compound  are  found  to  contain  16  parts  of  oxygen  and  28  parts  of 
nitrogen. 

The  molecular  weight  of  nitric  oxide  is  30 ;  30  parts  of  this 
compound  contain  16  parts  of  oxygen  and  14  parts  of  nitrogen. 

The  molecular  weight  of  nitrogen  is  found  to  be  28. 

In  these  four  different  molecules  the  weights  of  nitrogen  relative 
to  the  weight  of  i  atom  of  hydrogen  are  14,  28,  14,  28,  the 
greatest  common  divisor  of  which  is  14.  The  atomic  weight  of 
nitrogen,  therefore,  is  regarded  as  14. 

These  three  examples,  namely,  chlorine,  oxygen,  and  nitrogen 
are  instances  of  elements  which  are  gaseous  at  ordinary  tempera- 
tures ;  but  the  same  methods  are  applicable  in  the  case  of  the  non- 
volatile elements,  such  as  carbon,  provided  they  furnish  a  number 
of  compounds  that  are  readily  volatile. 

On  comparing  the  numbers  in  the  foregoing  table  (p.  42), 
representing  the  densities  of  various  elements,  with  the  atomic 
weights  of  those  elements  as  given  on  p.  22,  it  will  be  seen 
that  in  several  cases  the  numbers  given  are  approximately 
the  same.  This  agreement  is  merely  because  the  molecules 
of  these  elements  consist  of  two  atoms.  The  molecules  of 


44  Introductory  Outlines 

helium,  neon,  argon,  krypton,  xenon,  sodium,  potassium,  zinc, 
cadmium,  and  mercury  consist  of  only  one  atom  ;  their  atomic 
weights,  therefore,  will  be  the  same  as  their  molecular  weights,  that 
is,  twice  their  densities.  The  elements  arsenic  and  phosphorus,  on 
the  other  hand,  contain  in  their  molecules  four  atoms — that  is  to 
say,  the  number  which  represents  the  smallest  weight  of  phosphorus 
and  of  arsenic,  capable  of  taking  part  in  a  chemical  change,  is  only 
half  the  density,  and  therefore  a  fourth  of  the  molecular  weight. 

The  definition  of  atomic  weight  that  is  furnished  by  the  con- 
sideration of  volumetric  relations  may  be  thus  stated.  The  atomic 
weight  is  the  smallest  weight  of  an  element  that  is  ever  found  in  a 
volume  of  any  gas  or  vapour  equal  to  the  volume  occupied  by  one 
molecule  of  hydrogen  at  the  same  temperature  and  pressure. 

The  volume  occupied  by  one  molecule  of  hydrogen  is  regarded 
as  the  standard  molecular  volume,  while  that  occupied  by  an  atom 
of  hydrogen — or,  in  other  words,  the  atomic  volume  of  hydrogen — is 
called  the  unit  volume.  The  standard  molecular  volume,  therefore, 
is  said  to  be  two  unit  volumes;  and  as,  from  Avogadro's  law,  all 
gaseous  molecules  have  the  same  volume,  it  follows  that  the  mole- 
cules of  all  gases  and  vapours  occupy  two  unit  volumes.  Atomic 
weight  may  therefore  be  defined  as  the  smallest  weight  of  an 
element  ever  found  in  two  unit  volumes  of  any  gas  or  vapour. 

The  molecular  volume  of  a  gas  is  its  molecular  weight  divided 
by  its  relative  density,  a  ratio  which  in  all  cases  will  obviously 
equal  2,  that  is,  two  unit  volumes. 

The  atomic  volume  of  an  element  in  the  state  of  vapour  is  its 
atomic  weight  divided  by  its  relative  density.  In  the  case  of  such 
elements  as  chlorine,  nitrogen,  oxygen,  &c.,  whose  molecules  are 
diatomic,  the  quotient  will  be  i — that  is  to  say,  the  atomic  volume 
of  these  elements  is  equal  to  I  unit  volume.  In  the  case  of  mer- 

atomic  weight  =  200 

cury  vapour,  however,  we  have = : — ' =2. 

density  =100 

The  atomic  volume  of  mercury  vapour,  therefore,  is  equal  to  2 
unit  volumes,  and  is  identical  with  its  molecular  volume. 

On  the  other  hand,  with  the  element  phosphorus  the  atomic 

atomic  weight  =  31 
volume  is j —  .     °    —     =  -5j  or  one-half  the  unit  volume, 

and  therefore  one-fourth  the  molecular  volume  ;  consequently,  four 
atoms  exist  in  this  molecule. 

The  method  of  determining  atomic  weights  based  upon  volu- 
metric relations,  when  taken  by  itself,  is  not  an  absolutely  certain 


Atomic   Weights  45 

criterion,  for  although  the  atomic  weight  of  an  element  cannot  be 
greater  than  the  smallest  mass  that  enters  into  the  composition  of 
the  molecules  of  any  of  its  known  compounds,  it  might  be  less  than 
this,  as  there  is  always  the  possibility  of  a  new  compound  being 
discovered,  in  which  the  relative  weight  of  an  element  is  such  as  to 
make  it  necessary  to  halve  the  previously  accepted  atomic  weight. 
3.  Determination  of  Atomic  Weight  from  the  Specific 
Heat  of  Elements  in  the  Solid  State.— When  equal  weights  of 
different  substances  are  heated  through  the  same  range  of  tempera- 
ture, it  is  found  that  they  absorb  very  different  quantities  of  heat, 
and  on  again  cooling  to  the  original  temperature,  they  consequently 
give  out  different  amounts  of  heat.  Thus,  if  I  kilogramme  of  water, 
and  I  kilogramme  of  mercury  be  each  heated  to  a  temperature  of 
100°,  and  then  each  be  poured  into  a  separate  kilogramme  of  water 
at  o°,  in  the  first  case  the  resultant  mixture  will  have  a  temperature 
of  50°,  while  in  the  second  it  will  only  reach  the  temperature  of  3.2°; 
that  is  to  say,  while  the  water  in  cooling  through  50°  has  raised  the 
temperature  of  an  equal  weight  of  water  from  o°  to  50°,  the  amount 
of  heat  in  I  kilogramme  of  mercury  at  100°  has  only  raised  the 
temperature  of  an  equal  weight  of  water  from  o°  to  3.2°,  and  in  so 
doing  has  itself  become  lowered  in  temperature  100  -  3.2  =  96.8°.  The 
amount  of  heat  contained,  therefore,  in  equal  weights  of  water  and  of 
mercury  at  the  same  temperature,  as  shown  by  these  figures,  is  as — 

52.  M.-J.  i  . 
50  "96.8"    '^' 

therefore  it  requires  30  times  as  much  heat  to  raise  a  given  weight 
of  water  through  a  given  number  of  degrees  as  to  raise  an  equal 
weight  of  mercury  through  the  same  interval  of  temperature,  or 
the  thermal  capacity  of  mercury  is  3*0 th  that  of  water. 

The  specific  heat  of  a  substance  is  the  ratio  of  its  thermal 
capacity  to  that  of  an  equal  weight  of  water  ;  or,  the  ratio  between 
the  amount  of  heat  necessary  to  raise  a  unit  weight  of  the  sub- 
stance from  o°  to  i°,  and  that  required  to  raise  the  same  weight 
of  water  from  o°  to  i° ;  thus,  the  specific  heat  of  mercury  is  3^,  or 
0.033.  Water  is  chosen  as  the  standard  of  comparison  because  it 
possesses  the  highest  thermal  capacity  of  all  known  substances  ; 
the  numbers,  therefore,  which  express  the  specific  heats  of  other 
substances  are  all  less  than  unity. 

Dulong  and  Petit  were  the  first  to  draw  attention  (1819)  to  a 
remarkable  relation  which  exists  between  the  specific  heats  and 
the  atomic  weights  of  various  solid  elements,  whose  specific  heats 


46  Introductory  Outlines 

they  themselves  had  determined.  They  found  that  the  specific 
heats  of  the  solid  elements  were  inversely  as  their  atomic  weights ; 
that  is  to  say,  the  capacity  for  heat  of  masses  of  the  elements  pro- 
portional to  their  atomic  weight  was  equal.  This  law,  known  as 
the  law  of  Dulong  and  Petit,  may  be  thus  stated  :  The  thermal 
capacities  of  atoms  of  all  elements  in  the  solid  state  are  equal. 

The  thermal  capacity  of  an  atom  is  termed  its  atomic  heat; 
hence  the  law  may  be  more  briefly  stated,  all  elements  in  the 
solid  state  have  the  same  atomic  heat.  This  important  constant 
is  the  product  of  the  atomic  weight  into  the  specific  heat.  From 
the  following  table  it  will  be  seen  that  the  number  expressing 
the  atomic  heat  is  not  perfectly  constant :  the  departures  from  the 
mean .  6.4  are,  as  a  rule,  only  slight,  and  may  be  attributed  to 
the  fact  that  the  determinations  are  not  always  made  upon  the 
elements  under  conditions  that  are  strictly  comparable.  At  the 
end  of  the  table,  however,  there  are  certain  elements  which  appear 
to  present  marked  exceptions  to  the  law. 


Element. 

Specific 
Heat. 

Atomic 
Weight. 

/ 

] 

Ltomic 
Seat. 

Lithium    . 

0.94 

X 

7 

= 

6.6 

Sodium 

0.29 

X 

23 

= 

6.7 

Potassium 

o.i  66 

X 

39 

= 

6.5 

Manganese 

O.I  22 

X 

55 

= 

6.7 

0  112 

X 

56 

_ 

6.3 

Silver 

0.057 

X 

J 

1  08 

= 

J 

6.1 

Gold 

0.032 

X 

196 

= 

6.2 

Mercury  (solid) 

0.032 

X 

200 

= 

6.4 

Lead 

0.031 

X 

206.4 

= 

6.5 

/'Beryllium 

0.41 

X 

9-i 

= 

3-7 

1  Boron  (cryst.)  . 

0.25 

X 

ii 

** 

2.75 

j  Carbon  (diamond)    . 

0.147 

X 

12 

= 

1.76 

vSilicon  (cryst). 

0.177 

X 

28 

= 

4-95 

It  will  be  seen  that,  relatively  speaking,  the  four  elements 
which  show  a  considerable  departure  from  the  law  of  Dulong  are 
elements  with  low  atomic  weights.  Low  atomic  weight,  however, 
is  not  always  accompanied  by  such  deviation,  as  is  shown  in  the 
case  of  lithium  and  sodium. 

When  the  different  allotropes  of  carbon  are  experimented  upon, 
it  is  found  that  the  departure  is  not  the  same  for  each  modification 
of  the  element,  thus — 


Atomic    Weights  47 

•CM  Specific       Atomic      Atomic 

Heat.        Weight.       Heat. 

Diamond  .  .  .  0.147  x  12  =  1.76 
Graphite  .  .  .  0.200  x  12  =  2.40 
Charcoal  .  .  .  0.241  x  12  =  2.90 

It  has  been  observed  that,  as  a  general  rule,  the  specific  heat  of 
an  element  is  slightly  higher  at  higher  temperatures  ;  but  in  the 
case  of  the  four  elements  showing  abnormal  atomic  heats,  this 
increase  rises  rapidly  with  increased  temperature,  until  a  certain 
point  is  reached,  when  it  remains  practically  constant,  and  repre- 
sents an  atomic  heat  which  closely  approximates  to  the  normal 
value  ;  thus  in  the  case  of  diamond,  the  specific  heat  at  increasing 
temperatures  is — 

Specific         Atomic      Atomic 
Heat.          Weight.       Heat. 

Diamond  at  10.7°  .  .  .  0.1128  x  12  =  1.35 

„             45°  .  .  .  0.1470  x  12  =  1.76 

„           206°  .  .  .  0.2733  x  12  =  3.28 

„           607°  .  .  .  0.4408  x  12  =  5.30 

„           806°  .  .  .  0.4489  x  12  =  5.4 

985°        .  .  .      04589     X     12    =     5.5 

The  same  result  is  seen  in  the  case  of  graphite,  and  it  is  also  to 
be  remarked,  that  while  at  low  temperatures  there  exists  a  wide 
difference  between  the  specific  heats  of  these  two  modifications  of 
carbon,  this  difference  vanishes  at  a  temperature  of  about  600°. 

Specific         Atomic      Atomic 
Heat.          Weight.       Heat. 

Graphite  at  10.8°    .         .         .     0.1604   x    12   =    1.93 

61.3°    .         .        .     0.1990   X   12  —   2.39 

642°       .         .        .     04454   x    12   =   5.35 

„  978°          .  .  .      04670     X     12    =     5.50 

Both  the  elements  boron  and  silicon  are  found  to  follow  the 
same  rule,  and  at  moderate  temperatures  their  atomic  heats  nearly 
approximate  the  normal  constant. 

The  case  of  the  somewhat  rare  element  beryllium  is  of  special 
interest  from  another  point  of  view,  which  will  be  referred  to  when 
treating  of  the  natural  classification  of  the  elements  :  from  the 
following  numbers*  it  will  be  seen  that  its  atomic  heat  very 
rapidly  rises  with  moderate  increase  of  temperature. 
*  Humpidge. 


48  Introductory  Outlines 

Specific         Atomic     Atomic 
Heat.          Weight.       Heat. 

Beryllium  at  100°  .  .  .  0.4702  x  9.1  =  4.28 

„  200°  .  .  .  0.5420  x  9.1  =  4.93 

„  4°o°  •  •  •  0.6172  x  9.1  =  5.61 

„  500°  .  .  .  0.6206  x  9.  i  =  5.65 

The  relation  between  atomic  weight  and  specific  heat,  established 
by  Dulong  and  Petit,  is  of  service  in  the  determination  of  atomic 
weights,  not  as  a  method  of  ascertaining  the  exact  value  with  any 
degree  of  refinement,  but  rather  as  a  means  of  deciding  between 
two  numbers  which  are  multiples  of  a  common  factor. 

If  specific  heat  x  atomic  weight  =  atomic  heat,  it  will  be  obvious 
that,  if  we  experimentally  determine  the  specific  heat,  and  divide 
that  value  into  the  constant  atomic  heat,  6.4,  we  obtain  the 
approximate  atomic  weight. 

The  two  following  examples  will  serve  to  illustrate  the  applica- 
tion of  the  method. 

The  element  indium  combines  with  chlorine  in  the  proportion  — 

Indium  :  chlorine  =  37.8  :  35.5. 

If  InCl  is  the  formula,  then  37.8  is  the  atomic  weight  of  indium  ; 
but  from  the  chemical  similarity  between  indium  and  zinc  (whose 
chloride  has  the  formula  ZnCl2),  it  was  believed  that  the  formula 
for  indium  chloride  was  InCl2,  in  which  case,  in  order  to  preserve 
the  ratio  between  the  two  elements,  the  atomic  weight  would  have 
to  be  37.8  x  2  =  75.6. 

When  the  specific  heat  of  indium  was  determined,*  it  was  found 
to  be  0.057. 


Therefore  the  atomic  weight  must  be  raised  by  one-half,  from 
75.6  to  113.4,  and  the  formula  for  the  chloride  will  be  InCl3. 

The  element  thallium  combines  with  chlorine  in  the  proportion  — 

Thallium  :  chlorine  =  203.6  :  35.5. 

In  some  of  its  compounds  thallium  exhibits  a  strong  resemblance 
to  potassium,  the  chloride  of  which  has  the  formula  KC1.  If  the 
formula  for  the  thallium  chloride  is  T1C1,  the  atomic  weight  of  the 
metal  must  be  203.6. 

In  many  respects  thallium  exhibits  a  striking  analogy  with  lead, 
*  Bunsen,  1870. 


Atomic  Weights  49 

the  chloride  of  which  has  the  formula,  PbCl2.  If  thallium  chloride 
has  a  corresponding  formula,  T1C12,  then  the  atomic  weight  of 
thallium  must  be  raised  to  407.2. 

When  the  specific  heat  of  thallium  was  ascertained,*  it  was  found 
to  be  0.0335. 

6.4 


This  result  shows  that  the  number  203.6  and  not  407.2  is  the 
atomic  weight  of  thallium,  and  that  the  chloride  has  the  formula 
T1C1. 

Molecular  Heat  of  Compounds.—  The  capacity  for  heat  of  an 
atom  undergoes  no  alteration  when  the  atom  enters  into  combina- 
tion with  different  atoms  —  in  other  words,  the  atomic  heat  of  an 
element  is  the  same  in  its  compounds.  The  molecular  heat  of  a 
compound  (that  is,  the  product  of  the  molecular  weight  into  the 
specific  heat)  will  therefore  be  the  sum  of  the  atomic  heats  of  its 
constituent  elements.  Hence  it  is  possible  to  calculate  what  will 
be  the  atomic  heat  of  an  element  which  does  not  exist  as  a  solid 
under  ordinary  conditions  ;  and  therefore  the  atomic  weight  of 
such  an  element,  as  deduced  from  other  considerations,  is  capable 
of  verification,  by  determinations  of  the  molecular  heat  of  various 
of  its  compounds  :  thus  — 

The  specific  heat  of  silver  chloride,  AgCl,  is  0.089  :  — 

Specific  Molecular  Molecular 

Heat.  Weight.  Heat. 

0.089     x      143.5     =     I2'77- 

The  atomic  heat  of  silver  =  6.1,  therefore,  as  deduced  from  this 
compound,  the  atomic  heat  of  chlorine  is  12.77  —  6.1  =  6.6. 

Again,  the  specific  heat  of  stannous  chloride,  SnCl2,  is  0.1016:  — 

Specific  Molecular  Molecular 

Heat.  Weight.  Heat. 

0.1016      x      189     =       19.2. 

The  atomic  heat  of  tin  is  6.6,  therefore  the  atomic  heat  of  two 
atoms  of  chlorine,  as  deduced  from  this  compound,  is  19.2  —  6.6  = 
12.6,  giving  6.3  as  the  atomic  heat  of  chlorine. 

The  differences   that   appear  in   the   value,  as   deduced  from 

various   compounds,   are    lessened,   because    the    errors    of   the 

method  are  more  equally  distributed,  if  we  divide  the  molecular 

beat  by  the   number  of  atoms   in  the   molecule.     Thus,  in  the 

*  Regnault. 


50  Introductory  Outlines 

two  examples  quoted,  silver  chloride  consists  of  two  atoms,  while 
the  molecule  of  stannous  chloride  contains  three  ;  if,  therefore,  the 
molecular  heats  of  these  two  compounds  are  divided  respectively 
by  2  and  by  3  we  get  — 


as  the  value  representing  the  atomic  heat  of  chlorine. 

The  element  calcium  combines  with  chlorine  in  the  pioportion  — 

Calcium  :  chlorine  =  20  :  35.5. 

If  the  atomic  weight  of  calcium  is  20,  the  formula  will  be  CaCl, 
whereas  if  40  is  the  atomic  weight  of  the  metal,  the  compound 
must  be  represented  by  the  formula  CaCl2. 

The  molecular  weight  of  CaCl  would  be  55.5,  that  of  CaCU  1  1  i.o. 

When  the  specific  heat  of  the  compound  was  determined,  it 
was  found  to  be  0.1642.  In  order,  therefore,  to  decide  between 
the  two  values  for  the  atomic  weight  of  calcium,  we  calculate  the 
molecular  heat  from  both  of  the  molecular  weights,  and  divide  the 
result  by  the  number  of  atoms  in  the  molecule  in  each  case. 

On  the  supposition  that  Ca  =  2o,  and  that  CaCl  represents  the 
chloride  :  — 

Cad.        . 


Or,  if  Ca  =  4o,  and  CaCl2  is  the  formula  for  the  chloride,  then  — 

~  ,-,,  o.i642X  1  1  i.o 

CaCl.         .        .   -     3  -    _ 


The  number  6.07,  which  nearly  agrees  with  the  constant  6.4, 
decides  the  value  40  as  the  atomic  weight  of  calcium.  The 
element  calcium  is  one  of  those  metals  which  it  is  very  difficult  to 
isolate  and  obtain  in  a  state  of  purity,  but  when  in  recent  years 
the  specific  heat  of  this  metal  was  experimentally  determined,* 
it  was  found  to  be  0.1704  :  — 

0.1704x40  =  6.8. 

Thus  affording  direct  confirmation  of  the  value  40  for  the  atomic 
weight  of  calcium,  which  had  been  deduced  from  the  molecular 
heat  of  its  compounds. 

*  Bun  sen. 


Atomic  Weights  51 

Deductions  based  upon  molecular  heats  of  compounds  are  only 
trustworthy  in  the  case  of  the  most  simply  constituted  compounds. 

4.  Determination  of  Atomic  Weight  from  Considerations 
based  on  Isomorphism. — It  was  early  observed  that  certain  rela- 
tions existed  between  the  crystalline  forms  of  compounds  and  their 
chemical  composition.  Mitscherlich  found  that  certain  substances 
having  an  analogous  chemical  composition,  as,  for  example,  sodium 
phosphate  and  sodium  arsenate,  crystallised  in  the  same  geometric 
form.  In  the  year  1821  he  stated  his  law  of  isomorphism  as  follows  : 
"  The  same  number  of  atoms,  combined  in  the  same  way,  give  rise 
to  the  same  crystalline  form,  which  is  independent  of  the  chemical 
nature  of  the  atoms,  being  influenced  only  by  their  number  and 
mode  of  arrangement."  Subsequent  investigations,  however,  have 
shown  that  this  statement  is  too  general. 

In  its  broad  sense  as  signifying  the  same  crystalline  form, 
isomorphism  is  found  to  exist — 

1.  Between  compounds  containing  the  same  number  of  atoms 
similarly  combined,  and  which  bear  close  chemical  analogies  to 
each  other. 

c  (  Zinc  sulphate      ....     ZnSO4,7H2O. 

Isomorphous  <    .  ,  ,  , ,  _  *  __.f  _ 

I  Magnesium  sulphate  .         .         .     MgSO4,7H.2O. 

,          (  Hydrogen  disodium  phosphate  .     HNa2PO4,12H2O. 
13  (  Hydrogen  disodium  arsenate     .     HNa2AsO4,12H2O. 

/  Rubidium  alum  .         .         .         .     Rb2SO4,Al2(SO4)3>24H2O. 

Isomorphous  <  Potass!um  chrome  alum     '.       '     K2SO4,Cr2(SO4)3,24H2O. 
)  Potassium  aluminium  selenium  ) 

\       alum  ) 

2.  Between  compounds  containing  a  different  number  of  atoms, 
but  which  also  bear  close  chemical  analogies  to  one  another. 

,          (  Ammonium  chloride .         .         .     NH4C1. 
Isomorphous  < 

I  Potassium  chloride    .         .         .     KC1. 

Isomorphous  {  Ammonium  sulphate          .         .     (NH4)2SO4. 
(  Potassium  sulphate  .         .         .     K2SO4. 

3.  Between  compounds  containing  either  the  same  or  a  different 
number  of  atoms,  and  which  exhibit  little  or  no  chemical  analogies. 

T  ,          (  Sodium  nitrate  ....     NaNO3. 

Isomorphous  <  _  .  .  _  ^_  J 

(.  Calcium  carbonate    .        .         .     CaCO3. 

Isomorphous  {  Sodium  sulphate  (anhydrous)   .     Na2SO4. 
I  Barium  permanganate      .         .     BaMn2Ofl. 


52  Introductory  Outlines 

Isomorphism  of  this  order,  where  little  or  no  chemical  relations 
exist  between  the  compounds,  is  sometimes  distinguished  as 
isogonism.  It  must  not  be  supposed,  that  because  two  chemically 
analogous  compounds  contain  the  same  number  of  atoms,  they  will 
necessarily  crystallise  in  the  same  form  :  there  are  indeed  a  large 
number  of  similarly  constituted  analogous  compounds  that  do  not 
exhibit  isomorphism. 

No  simple  definition  of  isomorphism  is  possible,  but  the  following 
test  is  generally  accepted  as  a  criterion,  namely,  the  power  to  form 
either  mixed  crystals  or  layer  crystals.  Thus,  when  two  substances 
are  mixed  in  a  state  of  liquidity,  and  allowed  to  crystallise,  if  the 
crystals  are  perfectly  homogeneous,  they  are  known  as  mixed 
crystals,  and  the  substances  are  regarded  as  isomorphous. 

Or  when  a  crystal  of  one  compound  is  placed  in  a  solution  of 
another  compound,  and  the  crystal  continues  to  grow  regularly 
in  the  liquid,  the  compounds  are  isomorphous.  Thus,  if  a  crystal 
of  potassium  alum  (white)  be  placed  in  a  solution  of  manganese 
alum,  the  crystal  continues  to  grow  without  change  of  form,  and 
a  layer  of  amethyst-coloured  manganese  alum  is  deposited  upon  it. 

In  making  use  of  the  law  of  isomorphism  in  the  determination  of 
atomic  weights,  it  is  assumed  that  the  weights  of  different  atoms 
that  can  mutually  replace  each  other  without  altering  the  crystal- 
line form  are  proportional  to  their  atomic  weights.* 

Thus,  if  we  suppose  that,  in  the  case  of  the  sulphates  of  zinc 
and  magnesium,  the  atomic  weight  of  zinc  is  known,  viz.,  65,  and 
that  of  magnesium  is  doubtful  ;  from  the  fact  of  the  isomorphism 
of  the  sulphates  it  may  be  premised  that  the  elements  are  present  in 
proportions  relative  to  their  atomic  weights.  Analysis  shows  that 
the  proportion  is  24  of  magnesium  to  65  of  zinc,  therefore  24  is  pre- 
sumably the  atomic  weight  of  magnesium. 

In  this  way  Berzelius  corrected  many  of  the  atomic  weights 
which  in  his  day  had  been  assigned  to  the  elements. 

*  The  group  (NH4)  may  be  regarded  as  an  atom,  having  the  relative  weight  18. 


CHAPTER  VII 
QUANTITATIVE   CHEMICAL  NOTATION 

THE  use  of  chemical  symbols  and  formulae,  as  a  convenient  means 
of  representing  concisely  the  qualitative  nature  of  chemical  changes, 
has  been  explained  in  chapter  iv.  We  are  now  in  a  position  to 
read  into  these  symbols  a  quantitative  significance,  which  at  that 
stage  it  would  have  been  premature  to  explain. 

The  symbol  of  an  element  stands  for  an  atom  ;  but,  as  we  have 
now  learnt,  the  atoms  of  the  various  elements  have  different  relative 
weights,  hence  these  symbols  represent  relative  weights  of  matter. 
The  symbol  Na  signifies  23  relative  parts  by  weight  of  sodium,  O 
stands  for  16  relative  parts  by  weight  of  oxygen,  H  for  i  part  of 
hydrogen  ;  in  other  words,  the  weight  of  sodium  represented  by 
the  symbol  Na  is  23  times  as  heavy  as  that  which  is  conveyed 
by  a  symbol  H.  A  chemical  equation,  therefore,  is  a  strictly 
quantitative  expression,  in  which  certain  definite  weights  of  matter 
are  present  in  the  form  of  the  reacting  substances,  and  which 
reappear  without  loss  or  gain  in  the  compounds  resulting  from  the 
change.  In  this  sense  a  chemical  equation  is  a  mathematical 
expression.  Thus,  the  equation — 

Na  +  Cl  =  NaCl, 

not  only  means  that  an  atom  of  sodium  combines  with  an  atom  of 
chlorine  and  forms  i  molecule  of  sodium  chloride,  but  it  also  means 

23  +  35-5  =  58.5 
Na      Cl      NaCl. 

In  other  words,  that  sodium  and  chlorine  unite  in  the  relative  pro- 
portion of  23  parts  of  the  former  and  35.5  parts  of  chlorine,  and 
produce  58.5  parts  of  sodium  chloride. 

In  the  same  way,  into  the  equation  which  expresses  the  action  of 

53 


54  Introductory  Outlines 

sulphuric  acid  upon  sodium  carbonate,  we  read  the  quantitative 
meaning  of  the  symbols — 

H2SO4  +  Na2CO3  =  Na2SO4  +  CO2  +  H2O. 
2  46  46 

32  12  32  12  2 

64  48  64  32          16 

98*     +     106     =       142     +     44    +    1 8 

That  is  to  say,  98  parts  by  weight  of  sulphuric  acid  act  upon 
106  parts  of  sodium  carbonate,  producing  142  parts  of  sodium 
sulphate,  44  parts  of  carbon  dioxide,  and  18  parts  of  water.  It  will 
be  evident  that  it  becomes  a  matter  of  the  simplest  arithmetic  to 
calculate  the  weight  of  any  product  that  can  be  obtained  from  a 
given  weight  of  the  reacting  substances  ;  or  vice  versa,  to  find 
the  weight  of  any  reacting  substance  which  would  be  required  to 
produce  a  given  weight  of  the  product  of  the  action. 

Not  only  is  information  respecting  the  quantitative  relations 
by  weight  embodied  in  a  chemical  equation,  but  when  gaseous 
substances  are  reacting,  the  equation  also  represents  the  volu- 
metric relation  between  the  gases.  In  order  that  the  volumetric 
relations  may  be  more  manifest,  the  equations  expressing  the  re- 
actions are  written  in  such  a  manner  as  to  represent  the  molecules 
of  the  substances. 

H  +  C1  =  HC1 

is  an  atomic  equation,  but  as  the  molecule  is  the  smallest  particle 
which  can  exist  alone,  a  more  exact  statement  of  the  chemical 
change  is  made,  by  representing  the  action  as  taking  place  between 
molecules,  thus — 

H2  +  C12  =  2HCI. 

From  such  an  equation  we  see  that  i  molecule  of  hydrogen,  or 
2  unit  volumes,  unites  with  i  molecule  or  2  unit  volumes  of  chlorine, 
and  forms  2  molecules  or  4  unit  volumes  of  hydrochloric  acid  : 
or  again — 

02  +  2H2  =  2H20. 

One  molecule,  or  2  unit  volumes  of  oxygen,  unite  with  2  mole- 
cules, or  4  unit  volumes  of  hydrogen,  and  produce  2  molecules  of 

*  The  number  obtained  by  adding  together  the  weights  of  the  atoms  in  a 
formula  is  known  as  a  "  formula  weight,"  thus  98  is  the  formula  weight  of 
sulphuric  acid. 


Quantitative  Notanon  55 

water,  which  when  vaporised,  and  measured  under  the  same  con- 
ditions of  temperature  and  pressure,  occupy  4  unit  volumes.  In 
other  words,  the  number  of  molecules,  in  all  cases  *  where  gases 
and  vapours  are  concerned,  represent  exactly  the  volumetric 
relations.  In  the  cases  quoted,  it  will  be  observed,  the  same  ratio 
also  subsists  between  the  number  of  atoms  of  the*  reacting  gases 
and  the  molecules  of  the  compound,  but  this  is  not  always  the 
case,  for  example — 

Atomic  equation,  Hg  +  2C1  =  HgCl2. 

In  this  equation  3  atoms  unite  to  produce  i  molecule,  but  the 
ratio  between  the  volumes  is  not  represented  by  the  statement, 

1  volume  of  mercury  vapour  and  2  volumes  of  chlorine  produce 

2  volumes  of  vapour  of  mercury  chloride. 

Molecular  equation,  Hg  +  C12  =  HgCl2. 

By  this  we  see  that  i  molecule  t  (2  unit  volumes)  of  mercury 
vapour  and  i  molecule  (2  unit  volumes)  of  chlorine  give  i  mole- 
cule (2  unit  volumes)  of  vapour  of  mercury  chloride. 

Again, 

P  +  3C1  =  PC13 

is  an  atomic  equation,  showing  that  i  atom  of  phosphorus  unites 
with  3  atoms  of  chlorine  ;  but  it  is  not  true  that  the  ratio  between 
the  volumes  is  represented  by  the  statement,  I  volume  of  phos- 
phorus vapour  combines  with  3  volumes  of  chlorine  and  gives  2 
volumes  of  the  vapour  of  phosphorus  trichloride,  as  will  be  seen 
by  comparison  with  the  molecular  formulae — 

P4  +  6C12  =  4PC13. 

This  equation  tells  us  that  i  molecule  %  (2  unit  volumes)  of  phos- 
phorus vapour  combines  with  6  molecules  (12  unit  volumes)  of 
chlorine,  producing  4  molecules  (8  unit  volumes)  of  phosphorus 
trichloride  vapour.  / 

Knowing  the  relative  densities  of  gases  compared  with  hydro- 
gen, it  is  obviously  possible,  by  ascertaining  the  actual  weight  in 
grammes  of  some  definite  volume  of  hydrogen,  to  calculate  the 
actual  weight  of  any  given  volume  of  any  other  gas. 

Two  units  are  in  common  use,  namely — 

*  See  Dissociation,  where  apparent  exceptions  are  explained. 

f  The  atomic  volume  of  mercury  vapour  being  equal  to  2  unit  volumes  (p.  44). 

J  The  atomic  volume  of  phosphorus  is  .  5  of  a  unit  volume  (p.  44). 


56  Introductory  Outlines 

(i.)  The  weight  of  I  litre  of  hydrogen,  measured  at  a  temperature 
of  o°  C.,  and  under  a  pressure  of  760  mm.  of  mercury.* 

(2.)  The  volume  occupied  by  I  gramme  of  hydrogen,  measured 
under  the  same  conditions. 

I.  One  litre  of  hydrogen,  measured  at  the  standard  temperature 
and  pressure,  weighs  .0896  grammes.t  This  number  is  known  as 
the  crith;\  and  by  means  of  it  the  weight  of  I  litre,  and  therefore 
any  given  volume,  of  any  gas  can  be  deduced  :  thus,  the  relative 
densities  of  oxygen,  nitrogen,  and  chlorine  are  16,  14,  and  35.5 
respectively,  therefore  I  litre  of  these  gases  (measured  always  at 
the  standard  temperature  and  pressure)  weighs  16  criths,  14  criths, 
and  35.5  criths  respectively,  or  — 

i  litre  of  oxygen  weighs  16    x.o896=  1.4336  grammes. 
i       „       nitrogen     „      14     x.  0896=  1.2544        „ 
I       „       chlorine      „      35.5  x.  0896  =  3.  1808         „ 

So  also  with  reference  to  compound  gases,  where  in  each  case 
the  density  is  represented  by  the  half  of  the  molecular  weight. 
Thus,  the  relative  densities  of  hydrochloric  acid,  ammonia,  and 
carbon  dioxide  are  — 


and  the  weights  of  i  litre  of  these  gases  are  therefore  — 

I  litre  of  hydrochloric  acid=  18.25  x  .0896=  1.6352  gramme. 
i       „       ammonia  =  8.5    x.  0896  =  0.  7610        „ 

i       „       carbon  dioxide      =22.0   x.  0896  =1.97  12         „ 

II.  The  volume  occupied  by  i  gramme  of  hydrogen  at  the 
standard  temperature  and  pressure  is  11.127  litres.  As  the  rela- 
tive density  of  oxygen  is  16,  it  obviously  follows  that  16  grammes 

*  This  temperature  and  pressure  is  chosen  as  the  standard  at  which  volumes 
of  gases  are  compared.  See  General  Properties  of  Gases,  chapter  ix. 

f  From  time  to  time  slightly  different  values  have  been  given  for  this 
constant.  The  most  recent  determinations  give  the  number  .089873. 

£  From  the  Greek,  signifying  a  barley-corn,  and  used  symbolically  to  denote 
a  little  weight. 


Quantitative  Notation  57 

of  this  gas  will  also  occupy  11.127  litres;  in  other  words,  this 
number  11.127  represents  the  volume  in  litres  of  any  gas,  which 
will  be  occupied  by  the  number  of  grammes  corresponding  to  its 
relative  density,  thus — 

1 4  grammes  of  nitrogen    .         .     occupy  11.127  litres. 
35.5          „          chlorine     .         .          „        11.127      „ 
18.25        »          hydrochloric  acid        „       11.127      ?> 
22.0          „          carbon  dioxide .          „        11.127      ?? 

The  number  of  grammes  of  a  substance,  equal  to  the  number 
which  represents  its  molecular  weight,  is  spoken  of  as  the  gramme- 
molecule.  The  molecular  weight  of  hydrogen  =  2,  therefore  the 
gramme-molecule  of  hydrogen  (that  is,  2  grammes  of  hydrogen) 
will  occupy  11.127x2  =  22.25  litres.  The  molecular  weight  of 
oxygen  =  32,  therefore  32  grammes  of  oxygen  will  occupy  22.25 
litres  ;  in  other  words,  22.25  litres  is  the  volume  which  will  be 
occupied  by  the  gramme-molecule  of  any  gas. 

By  means  of  this  important  constant,  22.25,  tne  volume  of  any, 
or  all,  of  the  gaseous  products  of  a  chemical  change  (when 
measured  at  the  standard  temperature  and  pressure)  can  be  de- 
duced directly  from  the  equation  representing  the  change,  thus — 


expresses  the  reaction  taking  place  when  zinc  is  dissolved  in 
sulphuric  acid.  Just  as  in  the  former  illustrations  it  carries  the 
information  that  65  grammes  of  zinc +  98  grammes  of  sulphuric 
acid  produce  161  grammes  of  zinc  sulphate  and  2  grammes  of 
hydrogen.  But  2  grammes  of  hydrogen  occupy  22.25  litres,  there- 
fore by  the  solution  of  65  grammes  of  zinc,  the  volume  of  hydrogen 
obtained  will  be  22.25  litres. 

So  also  in  the  following  equation,  which  represents  the  formation 
of  carbon  dioxide  from  chalk  (calcium  carbonate)  by  the  action 
upon  it  of  hydrochloric  acid— 

CaCO3      +      2HC1      =      CaCl2      +       H2O      +       CO2. 

40+12  +  48       2(1+35.5)          40  +  71  2  +  16  12  +  32 

100         +         73         =        in         +        18         +         44 

100  grammes  of  chalk,  when  acted  upon  by  73  grammes  of  hydro- 
chloric acid,  yield  in  grammes  of  calcium  chloride  and  18 
grammes  of  water,  and  44  grammes  of  carbon  dioxide. 

Carbon  dioxide  is  gaseous,  therefore  44  grammes  (the  gramme- 


58  Introductory  Outlines 

molecule)  will  occupy,  at  the  standard  temperature  and  pressure, 
22.25  litres  ;  hence,  by  the  decomposition  of  100  grammes  of 
chalk,  22.25  h'tres  of  carbon  dioxide  are  produced. 

This  chapter  may  be  concluded  with  one  illustration  of  the 
methods  employed  in  the  exact  determination  of  atomic  weights 
which  depends  essentially  upon  the  quantitative  character  of 
chemical  reactions.  By  the  three  following  processes  the  atomic 
weights  of  chlorine,  potassium,  and  silver  may  be  deduced. 

1.  By  heating  a  known  weight  of  potassium  chlorate,  the  formula 
weight  of  potassium  chloride  is  found — 

KC1O3  =  KC1  +  3O. 

50  grammes  of  potassium  chlorate  when  heated  left  a  residue 
of  potassium  chloride  weighing  30.395  grammes.  50  —  30.395  = 
19.605  =  grammes  of  oxygen  evolved. 

As  potassium  chlorate  contains  in  its  formula  weight  3  atoms 
of  oxygen  (16  X  3  =  48),  we  get  the  expression — 

19.605  :  30.395=48  :  74.4o  =  formula  weight  of  potassium  chloride. 

2.  By  dissolving  a  known  weight  of  potassium  chloride,  and 
adding  to  it  excess  of  silve^-.nitrate,  silver  chloride  is  precipitated, 
which  can  be  washed  and  dried  and  weighed,  and  from  which  the 
formula  weight  of  silver  chloride  is  obtained — 

KC1  +  AgNO3  =  AgCl  +  KNO3. 

10  grammes  of  potassium  chloride  were  found  to  yield  19.225 
grammes  of  silver  chloride  ;  therefore/ 

10  :  19.225  =  74.40  :  143.03  =  formula  weight  of  silver  chloride. 

3.  By  the  direct  combination  of  silver  and  chlorine,  by  heating 
the  metal  in  a  stream  of  the  gas,  the  ratio  of  chlorine  to  silver  in 
silver  chloride  is  found  : 

10  grammes  of  silver  so  treated  yielded  13.285  grammes  of  silver 
chloride  ;  therefore, 

13.285  :  10  =  143.03  :  107.66  =  atomic  weight  of  silver. 
Since  the  formula  weight  of  silver  chloride,  AgCl  =  143.03, 

therefore,  143.03—  107.66  =  35.37  =  atomic  weight  of  chlorine. 
And  since  the  formula  weight  of  potassium  chloride,  KC1  =  74.40, 

therefore,  74.40  —  35.37  =  39.03  =  atomic  weight  of  potassium. 


CHAPTER  VIII 
VALENCY   OF   THE   ELEMENTS 

WHEN  chlorine  unites  with  hydrogen,  the  combination  takes  place 
between  one  atom  of  chlorine  (relative  weight  =  35.5)  and  one 
atom  of  hydrogen  (relative  weight  =  i) ;  but  when  oxygen  com- 
bines with  hydrogen,  one  atom  of  oxygen  unites  with  two  atoms 
of  hydrogen.  The  compound  ammonia  consists  of  one  atom  of 
nitrogen,  combined  with  three  atoms  of  hydrogen  ;  while  one  atom 
of  carbon,  on  the  other  hand,  can  unite  with  four  atoms  of 
hydrogen. 

One  atom  of  chlorine  never  combines  with  more  than  one  atom 
of  hydrogen  ;  its  affinity  for  that  element  is  satisfied,  or  saturated, 
by  union  with  one  atom. 

The  affinity  of  one  atom  of  oxygen  for  hydrogen,  however,  is 
not  satisfied  by  one  atom  of  that  element,  but  requires  two  atoms 
for  its  saturation  ;  while  nitrogen  requires  three,  and  carbon  four 
hydrogen  atoms,  in  order  to  satisfy  their  respective  affinities  for 
this  element. 

This  varying  power  of  combining  with  hydrogen  is  seen  in  a 
number  of  other  instances  :  thus,  the  elements  fluorine,  bromine, 
and  iodine,  resemble  chlorine  in  being  only  able  to  unite  with  one 
atom  of  hydrogen.  Sulphur,  like  oxygen,  has  its  affinity  for 
hydrogen  saturated  by  two  atoms  of  that  element.  Phosphorus 
and  arsenic  require  three  atoms  of  hydrogen  in  order  to  saturate 
their  combining  capacity,  while  silicon  resembles  carbon  in  com- 
bining with  four  hydrogen  atoms.  This  combining  capacity  of 
an  element  is  termed  its  valency.  Elements  like  chlorine, 
fluorine,  bromine,  and  iodine,  whose  atoms  are  only  capable 
of  uniting  with  one  atom  of  hydrogen,  are  called  monovalent 
(or  sometimes  monad)  elements  ;  while  those  whose  atoms  com- 
bine with  two,  three,  or  four  hydrogen  atoms,  are  distinguished 
as  di-valent  (or  dyad),  tri-valent  (or  triad),  and  tetra-valent  (or 
tetrad)  elements.  All  elements,  however,  are  not  capable  of 

59 


60  Introductory  Outlines 

entering  into  combination  with  hydrogen  ;  in  which  case,  their 
valency  is  measured  by  the  number  of  atoms  of  some  other 
monovalent  element  which  is  capable  of  satisfying  their  com- 
bining capacity.  Thus  : — 

atom  of  sodium  combines  with  i  atom  of  chlorine,  forming  NaCl. 
calcium  2  atoms  CaCL. 


boron          ,, 
tin  ,, 

phosphorus 
tungsten     , , 


BC13. 

SnCl4. 

PCL. 


WC16. 


In  the  combinations  of  elements  with  hydrogen  alone,  no  in- 
stances are  known  in  which  a  higher  valency  is  exhibited  than 
that  of  four  ;  but  with  chlorine,  as  here  seen,  cases  are  known  in 
which  elements  exhibit  pentavalent  and  hexavalent  characters. 

Measured  by  their  combining  capacity  for  hydrogen  and  chlorine, 
elements  do  not,  however,  always  exhibit  the  same  valency  :  thus, 
the  affinity  of  phosphorus  for  hydrogen  is  satisfied  by  three  hydrogen 
atoms,  whereas  one  atom  of  this  element  can  unite  with  five  atoms 
of  chlorine. 

As  measured  by  hydrogen,  the  valency  of  sulphur  is  two,  the 
compound  that  it  forms  with  hydrogen  being  expressed  by  the 
formula  SH2,  while,  as  estimated  by  its  capacity  for  chlorine,  it 
becomes  tetravalent,  as  seen  in  the  compound  SC14.  As  a  general 
rule,  however,  the  highest  number  of  monovalent  atoms  with  which 
one  atom  of  an  element  is  capable  of  combining  is  accepted  as 
representing  the  valency  of  that  element.  Thus,  one  atom  of 
phosphorus  not  only  combines  with  five  atoms  of  chlorine,  but 
also  with  five  atoms  of  fluorine  ;  phosphorus  is  therefore  a  penta- 
valent element. 

As  measured  by  hydrogen  alone,  or  by  chlorine  alone,  nitrogen 
is  a  trivalent  element,  for  the  largest  number  of  these  atoms  with 
which  one  atom  of  nitrogen  can  unite  is  three,  as  seen  in  the 
compounds  having  the  composition  NH3  and  NC13  ;  neverthe- 
less, one  atom  of  nitrogen  is  capable  of  combining  with  four 
atoms  of  hydrogen  and  one  of  chlorine,  forming  the  compound 
NH4C1,  ammonium  chloride,  in  which  the  nitrogen  atom  is  penta- 
valent. 

This  rule,  however,  is  not  always  followed  ;  for  example,  one 
atom  of  iodine  will  unite  with  three  atoms  of  chlorine,  forming  the 

*  Phosphorus  also  combines  with  hydrogen. 


Valency  of  the  Elements  61 

compound  IC13,  but  iodine  is  not  generally  regarded  as  a  trivalent 
element.* 

In  symbolic  notation,  this  power  possessed  by  an  atom,  of  uniting 
to  itself  monovalent  atoms,  is  often  represented  by  lines,  each  line 
signifying  the  power  of  combination  with  one  monovalent  atom. 
Thus,  in  the  symbol  H — Cl,  the  line  is  intended  to  give  a  concrete 
expression  to  the  fact  that  both  hydrogen  and  chlorine  are  mono- 
valent elements,  and  that  the  affinity  of  each  element  for  the 
other  is  satisfied  when  one  atom  of  the  one  unites  with  one  atom  of 
the  other.  The  symbol  H — O — H,  in  like  manner,  signifies  that 
the  oxygen  atom  is  divalent,  that  its  affinity  for  hydrogen  is  satisfied 
only  when  it  has  united  with  two  monad  atoms.  In  the  same  way 
we  may  express  the  facts  that  nitrogen  and  carbon,  in  their  com- 
binations with  hydrogen,  are  respectively  trivalent  and  tetravalent, 

H 

by  the  symbols  H — N — H,  and  H — C — H.     These  lines  are  merely 

H  H 

a  convenient  symbolic  expression  for  the  operation  of  the  force  of 
chemical  affinity  ;  their  length  and  direction  bear  no  meaning.t 
The  power  to  combine  with  one  monovalent  atom  is  sometimes 
spoken  of  simply  as  one  affinity  :  thus  it  is  said  that  in  the  com- 
pound having  the  composition  PH3,  or  H — P — H,  three  of  the 

H 

affinities  of  the  phosphorus  atom  are  saturated,  and  that  two 
affinities  still  remain  unsatisfied,  phosphorus,  as  already  stated, 
being  a  pentavalent  element. 

*  See  Iodine,  Compounds. 

f  The  student  cannot  be  too  often  warned  against  attaching  any  materialistic 
significance  to  these  lines.  The  use  of  this  convention  is  always  attended  with 
the  danger  that  the  beginner  is  liable  to  fall  into  the  error  of  regarding  these 
lines  as  representing  in  some  manner  fixed  points  of  attachment,  or  links, 
between  the  atoms.  It  must  be  remembered,  therefore,  that  these  lines  not  only 
have  no  materialistic  signification,  but  they  must  not  even  be  regarded  as  convey- 
ing any  statical  meaning.  The  atoms  are  undergoing  rapid  movements  with 
respect  to  each  other,  which  movements  are  in  some  way  governed  by  the 
chemically  attractive  force  exerted  by  the  individual  atoms  upon  one  another ; 
and  the  molecule  will  be  more  correctly  considered,  if  we  regard  its  atoms  as 
being  held  together  in  a  manner  resembling  that  by  which  the  numbers  of  a 
cosmical  system  are  bound  together.  The  lines  simply  denote  that  the  atoms 
are  held  to  each  other  by  the  attractive  force  which  we  call  chemical  affinity. 


62  Introductory  Outlines 

Compounds  of  this  order,  in  which  one  of  the  elements  has  still 
unsatisfied  affinities,  are  called  unsaturated  compounds. 

In  its  power  to  satisfy  the  affinities  of  an  element,  a  divalent 
atom  is  equal  to  two  monovalent  atoms  :  thus,  when  the  affinities 
of  the  tetravalent  carbon  atom  are  saturated  with  oxygen,  the  mole- 
cule contains  two  atoms  of  oxygen,  which  may  be  symbolically 
expressed  thus,  O  =  C  =  O,  in  which  the  four  affinities  of  the 
carbon  (represented  by  the  four  lines)  are  satisfied  by  the  two 
divalent  atoms  of  oxygen.  Carbon,  however,  combines  with  a 
smaller  proportion  of  oxygen,  forming  the  compound  carbon  mon- 
oxide, CO.  The  carbon  atom  in  this  case  is  divalent,  as  expressed 
by  the  formula  C  =  O,  and  this  substance  is  also  an  unsaturated 
compound. 

The  number  of  divalent  atoms  with  which  an  element  can  unite 
cannot,  however,  be  taken  as  a  safe  criterion  or  measure  of  the 
valency  of  that  element  in  cases  where  that  number  is  greater 
than  i  ;  for  example,  in  such  a  compound  as  calcium  oxide,  CaO, 
we  regard  the  two  affinities  of  the  divalent  atom  of  oxygen  as  being- 
satisfied  by  two  affinities  possessed  by  the  calcium,  and  express  this 
belief  in  the  formula  Ca  =  O,  and  regard  the  calcium  as  divalent. 
In  the  same  way,  in  carbon  monoxide,  CO,  the  carbon  being  united 
with  one  atom  of  the  divalent  element  oxygen  is  itself  divalent  in 
this  compound  ;  but  in  the  case  of  carbon  dioxide,  where  the  carbon 
atom  is  united  with  two  atoms  of  divalent  oxygen,  we  are  not 
justified  in  asserting  that  the  atoms  are  united,  as  represented  by 
the  formula  O  =  C  =  O,  in  which  the  four  affinities  of  carbon 
are  represented  as  saturated  with  oxygen.  There  exists  no  posi- 
tive proof  that  the  carbon  is  not  divalent  in  this  compound,  and 
that  the  molecule  does  not  consist  of  three  divalent  atoms  united, 

C 

as  shown  in  the  formula  /\ .      From  the  fact,  however,  that  car- 
O O 

bon  forms  a  compound  with  four  atoms  of  hydrogen,  and  another 
with  four  atoms  of  chlorine,  we  know  that  this  element  is  tetra- 
valent, and  therefore  we  believe  that  in  carbon  dioxide  it  is  also 
tetravalent. 

Again,  as  measured  by  its  compound  with  hydrogen,  sulphur  is 
divalent  ;  while  with  chlorine  it  forms  SC14.  But  sulphur  unites 
with  oxygen,  forming  the  two  compounds  sulphur  dioxide,  SO2,  and 
sulphur  trioxide,  SO3.  If  it  be  assumed  that  in  these  molecules  the 


Valency  of  the  Elements  63 

whole  of  the  oxygen  affinities  are  satisfied  with  sulphur,  then  the 
symbolic  representation  of  these  oxides  will  be  O  =  S  =  O,  and 
O  =  S  =  O,  the  sulphur  being  in  one  case  tetravalent  and  in  the 

II 
O 

other  hexavalent.  There  is,  however,  no  positive  proof  that  the 
affinities  of  one  oxygen  atom  are  not  partially  satisfied  by  union 
with  another  oxygen  atom,  and  that  the  valency  of  the  sulphur  is 
higher  than  either  two  or  four,  as  seen  in  the  alternative  formulas, 


S  ?\  S 

S02/\  S0:5  S  =  0;  or      /     \ 

O O  O— O— O 


Chemists  believe,  however,  that  in  these  two  oxides  the 
sulphur  functions  in  the  one  case  as  a  tetravalent,  and  in 
the  other  as  a  hexavalent  element;  and  this  belief  is  strengthened 
by  the  recent  discovery  (Moissan)  of  a  fluoride  having  the  com- 
position SF6,  in  which  the  hexavalent  character  of  sulphur  is 
unquestionable. 

It  will  be  evident  from  these  considerations,  that  in  many  cases 
the  valency  of  an  element  is  a  variable  quantity,  depending  partly 
upon  the  particular  atoms  with  which  it  unites.  It  is  also  found 
that  it  is  dependent  in  many  instances  upon  temperature  and 
upon  pressure.  Thus,  between  a  certain  limited  range  of 
temperature,  one  atom  of  phosphorus  combines  with  five  atoms 
of  chlorine  in  the  compound  PC15,  but  above  that  limit  two  atoms 
of  chlorine  leave  the  molecule,  and  the  phosphorus  becomes  tri- 
valent.  Again,  if  hydrogen  phosphide,  PH3,  be  mixed  with  hydro- 
chloric acid,  HC1,  and  the  mixed  gases  be  subjected  to  increased 
pressure,  the  gases  combine  and  form  a  solid  crystalline  com- 
pound known  as  phosphonium  chloride,  PH4C1,  in  which  the 
phosphorus  atom,  being  united  with  five  monovalent  atoms,  is 
pentavalent.  When  the  pressure  is  released  an  atom  of  hydrogen 
and  an  atom  of  chlorine  leave  the  molecule,  and  the  phosphorus 
returns  to  its  trivalent  condition. 

A  compound,  in  whose  molecules  there  is  an  atom  which  for  the 
time  being  is  not  functioning  in  its  highest  recognised  valency, 
often  exhibits  a  readiness  to  unite  with  additional  atoms  to  form 


64  Introductory   Outlines 

new    compounds  :  thus  ammonia   combines  eagerly  with  hydro- 
chloric acid,  forming  ammonium  chloride — 

NH3  +  HC1  =  NH4C1. 

Carbon  monoxide  unites  directly  with  chlorine  to  form  carbonyl 
chloride — 

=  COC1,. 


Carbon  monoxide  also  combines  with  an  additional  atom  of 
oxygen,  and  gives  carbon  dioxide,  thus — 

2CO  +  O2  =  2CO2. 

In  this  last  action  it  will  be  seen  that  the  molecule  of  carbon 
monoxide,  in  being  converted  into  the  dioxide,  takes  up  one  atom 
of  oxygen  ;  but  as  the  molecule  of  oxygen  is  the  smallest  isolated 
particle,  it  follows  that  the  two  atoms  contained  in  such  a  molecule 
must  first  separate,  and  each  one  then  furnishes  the  requisite 
additional  oxygen  for  one  molecule  of  carbon  monoxide.  In  the 
union  of  carbon  monoxide  with  chlorine,  and  of  ammonia  with 
hydrochloric  acid,  are  we  to  suppose  that  the  same  action  takes 
place  ?  That  is  to  say,  do  the  two  atoms  in  the  molecule  of 
chlorine  separate  from  each  other  and  unite  with  carbon,  thereby 
satisfying  its  tetrad  valency,  in  the  manner  here  expressed? — 

Ck 
Cl— — C1  +  CO=       >C  =  0. 

cy 

And  in  the  case  of  ammonia  and  hydrochloric  acid,  do  the 
hydrogen  and  chlorine  atoms  part,  and  each  unite  with  the 
nitrogen  atom,  thereby  raising  it  from  the  trivalent  to  the  penta- 
valent  condition  ?  thus — 

Cl        H 

;  \/ 

H— Cl  +  H— N— H  =  H— N— H. 

i  !  I 

H  H 


Valency  of  the  Elements  65 

Or  are  we  to  suppose  that  the  two  molecules,  without  losing  their 
integrity,  become  held  together  as  independent  molecules,  by 
virtue  of  the  unsatisfied  affinities  of  the  carbon,  or  the  nitrogen, 
as  the  case  may  be,  in  which  case  the  compounds  might  be  repre- 
sented thus — 

Cl  H  —  Cl 

|C  =  O  H  — N  — H. 

Cl  I 

H 

This  question  would  be  settled  by  determining  the  vapour- 
density  of  the  compound.  If,  for  instance,  we  were  to  find  the 
vapour-density  of  ammonium  chloride  to  be  26.75,  then  the  com- 
pound having  the  composition  NH^Cl  would  have  the  normal 
molecular  volume,  that  is,  its  molecule  would  occupy  two  unit 
volumes,*  and  the  conclusion  would  be  that  the  vapour  consisted 
of  single  molecules  of  the  composition  represented  by  the  formula 
NH4C1.  But  ammonium  chloride  at  ordinary  temperatures  is  a 
solid_,  and  when  heated  to  the  temperature  necessary  to  convert  it 
into  vapour  its  molecules  break  up  into  separated  molecules  of  the 
two  original  gases — ammonia,  NH3,  and  hydrochloric  acid,  HCl.t 
So  that  we  are  unable  to  gain  any  information  in  this  direction 
as  to  the  mode  in  which  the  atoms  are  disposed  in  the  compound. 
When  the  two  gases  are  brought  together  under  ordinary  con- 
ditions, they  combine  with  the  evolution  of  considerable  heat, 
owing  to  loss  of  energy  ;  this  is  taken  as  evidence  that  true 
chemical  action,  in  the  sense  of  atomic  rearrangement,  has  re- 
sulted, hence  it  is  believed  that  in  this  compound  the  nitrogen 
is  united  with  the  five  monovalent  atoms,  and  consequently  Is 
pentavalent. 

In  the  case  of  carbonyl  chloride,  COC12,  the  vapour-density  can 
be  ascertained,  this  compound  existing  in  the  gaseous  condition 
at  the  ordinary  temperature.  Its  vapour-density,  determined  by 
experiment,  is  found  to  be  50.6.  This  number,  divided  into  the 
molecular  weight  of  the  compound  having  the  composition  COCI2, 
gives  practically  the  number  2  as  the  molecular  volume  of  the 
compound.  Hence  we  conclude  that  these  four  atoms  constitute 
a  single  molecule. 

There  are  a  number  of  combinations,  however,  in  which  mole- 

*  See  p.  43.  f  See  Dissociation,  p.  89. 

E 


66  Introductory  Outlines 

cules  of  different  compounds  unite,  that  do  not  so  readily  admit  of 
explanation,  because  in  neither  of  the  molecules  is  there  any 
atom  functioning  in  a  lower  state  of  valency  than  that  which 
it  is  known  to  be  capable  of.  For  example,  the  monovalent 
elements  fluorine  and  hydrogen  form  the  compound  hydrofluoric 
acid,  HF  ;  fluorine  also  combines  with  the  monovalent  element 
potassium,  forming  potassium  fluoride,  KF.  Both  of  these  com- 
pounds come  under  the  head  of  saturated  compounds,  in  the  sense 
that  neither  of  them  contains  an  atom  which  is  known  to  be 
capable  of  exercising  a  higher  valency  than  it  exhibits  in  these 
compounds.  Nevertheless  these  two  molecules  unite  together  and 
form  a  definite  chemical  compound,  known  as  hydrogen-potassium 
fluoride. 

Again,  the  divalent  element  zinc  combines  with  two  atoms  of 
the  monad  element  chlorine,  forming  zinc  chloride,  ZnCl^ ;  the 
two  monovalent  elements  sodium  and  chlorine  also  combine, 
giving  the  compound  sodium  chloride,  NaCl.  Both  of  these 
substances  must  be  regarded  as  saturated  compounds,  and  yet 
they  unite  with  each  other,  forming  a  distinct  chemical  compound, 
known  as  sodium  zinc  chloride.  Such  compounds  as  these  are 
known  as  double  salts,  and  examples  might  be  multiplied  almost 
indefinitely.  A  similar  union  of  molecules,  where  the  recognised 
valency  of  the  atoms  is  all  satisfied,  is  seen  in  a  large  number 
of  compounds  containing  water  of  crystallisation ;  *  for  example, 
the  divalent  element  copper,  in  combination  with  two  atoms  of 
chlorine,  forms  cupric  chloride,  CuCl2.  The  divalent  element 
oxygen,,  in  combination  with  two  hydrogen  atoms,  forms  water, 
H2O.  When  cupric  chloride  crystallises  from  aqueous  solution, 
each  molecule  of  the  chloride  unites  to  itself  two  molecules  of 
water,  which  is  therefore  termed  water  of  crystallisation. 

In  chemical  notation,  it  is  usual  to  represent  compounds  of  this 
order  by  placing  the  formulas  of  the  different  molecules  that  have 
entered  into  union  in  juxtaposition,  with  a  comma  between ; 
accordingly,  the  examples  here  quoted  would  be  indicated  thus — 

Hydrogen  potassium  fluoride       .         .     HF,KF. 
Sodium  zinc  chloride   ....     ZnCl2,NaCl. 
Crystallised  cupric  chloride          .         .     CuCl2,2H2O. 

Combinations  of  this  order  are  by  no  means  confined  to   the 
*  See  page  216. 


Valency  of  the  Elements  67 

union  of  two  kinds  of  molecules,  as  the  following  examples  will 
serve  to  show  : — 

Platinum  sodium  chloride     .         .     PtCl4,2NaCl,6H,O. 
Mercuric  potassium  chloride          .     2HgCl2,KCl,2H2O. 

At  the  present  time  our  knowledge  of  the  nature  of  the  union 
between  these  various  molecules  is  too  imperfect  to  admit  of  any 
precise  explanation  ;  such  compounds  are  frequently  distinguished 
as  molecular  combinations. 

It  is  quite  possible  that  the  unit  which  has  been  adopted  for  estimating 
valency,  namely,  i  monovalent  atom,  is  after  all  only  an  extremely  rough  and 
crude  measure,  which  is  incapable  of  appreciating  smaller  differences  of  com- 
bining capacity  that  may,  and  most  probably  do,  exist.  Its  use  may  be  com- 
pared to  the  adoption  of  a  single  unit,  say  I  gramme,  for  the  estimation  of 
mass  or  weight ;  when,  if  a  given  quantity  of  matter  has  a  weight  equal  to  i 
gramme,  but  less  than  2  grammes,  its  weight  would  be  i;  if  greater  than  2 
grammes,  but  less  than  3,  then  its  weight  would  be  2 — a  method  of  estimating 
which  tacitly  assumes  that  no  intermediate  weights  of  matter  between  the 
various  multiples  of  the  selected  unit  are  possible.  There  is  no  evidence  to 
show  that  the  combining  capacity  of  an  element  is  exactly  expressed  by  simple 
multiples  of  a  monovalent  atom. 

For  example,  i  hydrogen  atom  unites  with  i  chlorine  atom,  that  is  to  say, 
with  a  mass  of  chlorine  weighing  35.5  times  its  own  weight ;  and  we  say  that 
the  mutual  affinities  of  these  atoms  are  satisfied.  But  for  anything  we  know 
to  the  contrary,  an  atom  of  hydrogen  may  have  an  affinity  for  chlorine  which 
would  enable  it  to  unite  with  a  mass  of  chlorine  weighing  40  or  45  or  50  times 
its  own  weight,  but  not  a  mass  weighing  71  (35.5  x  2)  times  its  own.  But  since 
a  mass  of  chlorine  35.5  times  the  weight  of  a  hydrogen  atom  is  the  smallest 
quantity  that  is  ever  known  to  take  part  in  a  chemical  change,  is  the  chemically 
indivisible  mass  we  call  an  atom,  it  follows  that  as  the  hydrogen  atom  has  not 
sufficient  combining  capacity  to  unite  with  2  atoms,  it  is  compelled  to  be 
satisfied  with  i.  It  might  still,  however,  retain  a  residual  combining  capacity, 
Or  the  residual  combining  capacity  may  be  lodged  in  the  chlorine  atom, 
which  may  be  conceived  as  being  able  to  unite  with  a  greater  weight  of 
hydrogen  than  is  represented  by  i  atom,  but  not  so  much  as  that  of  2 
atoms. 

Each  of  the  elements  fluorine,  chlorine,  bromine,  and  iodine  unites  with 
i  atom  of  hydrogen,  and  we  represent  their  compounds  in  a  similar  manner, 
thus — 

H  -  F  •         H  -  Cl ;         H  -  Br  ,         H  -  I ; 

but  we  make  an  enormous  assumption  if  we  suppose  that  in  each  of  these 
compounds  the  mutual  affinities  of  the  atoms  is  equally  satisfied. 

The  trend  of  modern  thought,  however,  lies  in  the  direction  of  an  electrical 
interpretation  of  valency.  The  fact  that  atoms  are  always  associated  with 
fixed  and  definite  charges  of  electricity,  that  valency ,  indeed,  could  be  measured 
in  terms  of  electric  units  (the  outcome  of  Faraday's  Law,  chap,  xi.)  seemed 


68  Introductory  Outlines 

at  one  time  only  to  emphasise  the  difficulty  of  explaining  such  cases  as  those 
above  mentioned  ;  but  the  more  recent  developments  in  this  region  of  physics 
have  led  to  modified  views  as  to  the  nature  of  the  bond  which  unites  atoms 
together.  Stated  in  briefest  outline,  this  chemical  "  bond"  or  unit  of  affinity, 
which  formerly  has  been  regarded  in  the  light  of  a  single  line  of  force — a 
fraction  of  a  bond  being  considered  as  altogether  inadmissible — is  now  regarded 
as  a  bundle  of  lines  of  force  (a  Faraday  bundle).  Under  appropriate  condi- 
tions, such  as  the  proximity  of  suitable  molecules  or  ions,  it  is  conceived  that 
some  strands  of  the  bundle  may  become  loosened  from  one  of  the  attached 
atoms  and  thus  become  available  for  attraction  by  similar  wandering  strands 
from  other  molecules.  Obviously,  therefore,  this  view  admits  of  practically  an 
unbroken  gradation  in  degrees  of  chemical  affinity.  Instead,  therefore,  of 
residual  affinity,  we  have  varying  fractions  of  the  total  bundle  of  lines  of 
force  which  in  its  entirety  constitutes  the  chemical  "  bond." 

A  modification  of  this  view,  recently  advanced  by  Sir  W.  Ramsay,*  substitutes 
electrons^  for  this  "  bundle  of  lines  offeree."  Atoms  are  regarded  as  carrying 
with  them  a  "  reserve  of  electrons  "  electrons  w:hich  may  be  inactive,  or  latent. 
Thus  taking  chlorine  as  an  example,  he  says:  "It  appears  likely  that  each 
atom  of  chlorine  carries  with  it  no  fewer  than  seven  electrons,  .  .  .  latent  as  it 
were,  not  revealing  themselves  in  such  a  compound  as  common  salt.  .  .  . 
These  valencies  are  manifested  in  such  compounds  as  perchloric  acid." 

*  Presidential  address,  Chem.  Soc.,  1909.  f  See  page  104. 


CHAPTER  IX 
GENERAL   PROPERTIES    OF   GASES 

UNDER  the  head  of  the  general  properties  of  gases  it  will  be  con- 
venient to  consider  the  following  subjects  :* — 

1.  The  relation  of  gases  to  heat. 

2.  The  relation  of  gases  to  pressure. 

3.  The  liquefaction  of  gases. 

4.  Diffusion  of  gases. 

5.  The  kinetic  theory  of  gases. 

The  Relation  of  Gases  to  Heat.— The  fact  that  substances 
expand  when  heated,  and  again  contract  upon  being  cooled,  was 
observed  in  very  early  times.  The  fact  also  that  all  substances  do 
not  undergo  the  same  alterations  in  volume  when  subjected  to  the 
same  changes  of  temperature  has  been  long  known  ;  but  it  was  not 
until  the  beginning  of  the  nineteenth  century  that  it  was  proved  by 
Charles  and  .Gay-Lussac  that  all  gases  expanded  and  contracted 
equally  when  exposed  to  the  same  alterations  of  temperature. 
This  law  is  generally  known  as  the  Law  of  Charles,  and  may  be 
thus  stated  :  When  a  gas  is  heated,  the  pressure  being  constant^  it 
increases  in  volume  to  tJie  same  extent  whatever  the  gas  may  be. 

The  increase  in  bulk  suffered  by  I  volume  of  a  gas  in  being 
heated  from  o°  to  i°  is  termed  the  coefficient  of  expansion,  and  if 
the  law  of  Charles  is  true  all  gases  will  have  the  same  coefficient. 

Modern  research  has  shown  that  the  law  of  Charles  is  not  abso- 
lutely true,  and  the  extent  to  which  gases  deviate  from  the  strict 
expression  will  be  seen  from  the  coefficients  of  expansion  given  in 
the  following  table  : — 

*  The  study  of  these  subjects  belongs  more  especially  to  the  science  of 
physics  or  chemico-physics.  For  fuller  information  on  these  points  than  can 
be  included  within  the  scope  of  this  book  students  are  referred  to  special 
treatises  on  physics. 


70  Introductory  Outlines 

Air 003665' 

Hydrogen 003667 

Carbon  monoxide 003667 

Nitrogen oo3668j 

Nitrous  oxide     ..,.„.     .003676 
Carbon  dioxide  ......     .003688 

Cyanogen 003829 

Sulphur  dioxide          ....     .003845 

It  will  be  noticed  that  the  first  four  gases  have  almost  the  same 
coefficient  of  expansion  :  these  gases  are  all  very  difficult  of  lique- 
faction, and  it  will  be  seen  that  the  coefficient  rapidly  rises  in  the 
case  of  the  other  gases,  which  are  easily  liquefied. 

For  purposes  of  ordinary  calculation  it  is  usual  to  adopt  the 
coefficient  of  expansion  of  air  as  applicable  to  all  gases.  It  will 
be  obvious  that  since  the  volume  of  a  gas  is  affected  by  alterations 
of  temperature,  it  becomes  necessary,  when  measuring  the  volume 
of  a  gas,  to  have  regard  to  the  particular  temperature  at  which  the 
measurement  is  made,  and  in  order  to  compare  volumetric  measures 
they  must  be  all  referred  to  some  standard  temperature.  This 
standard  temperature  is  by  general  consent  o°  C. 

Taking  the  fraction  .003665,  therefore,  for  the  coefficient — 

I  volume  of  a  gas  at  o°  becomes  I  +  .003665  volumes  at  i° 

I  „  „          o°         „         i  +  .003665  x  2     „         2° 

or       i  „          „          o°         „         i  +  .003665 1          „         t° 

Therefore  the  volume  at  /°  equals  the  volume  at  o°  multiplied  by 
i  +  .003665  /.  Let  "v  be  the  volume  at  /°,  and  v0  the  volume  at  o°, 
then — 

v  =  v0(i  +  .0036654 

and  conversely  the  volume  at  o°  equals  the  volume  at  t°  divided  by 
i  +  .003665  t— 


i  +  .003665  / 

The  vulgar  fraction  equivalent  to  .003665  is  of 5-.  273  volumes 
at  o°  become  273  -J-  /  at  f. 

What  is  known  as  the  absolute  temperature  of  a  substance  is  the 
number  of  degrees  above  -  273°  C.  Taking  this  point  as  the  zero, 
the  absolute  temperature  of  melting  ice,  for  example,  will  be  273°. 
Charles'  law,  therefore,  may  be  thus  stated  :  The  volume  of  any 


Relation  of  Gases  to  Pressure  71 

gas,  under  constant  pressure^  is  proportio7ial  to  the  absolute  tern* 
Perature. 

The  Relation  of  Gases  to  Pressure.—  The  effect  of  increase 
of  pressure  upon  a  gas  is  to  diminish  its  volume.  The  law  which 
connects  the  volume  occupied  by  a  gas,  with  the  pressure  to  which  it 
is  subjected,  was  discovered  by  Robert  Boyle  (1661),  and  is  known 
as  Boyle's  Law.  It  may  be  thus  stated  :  The  volume  occupied  by 
a  given  weight  of  any  gas  is  inversely  as  the  pressure.  The 
general  truth  of  this  law  may  readily  be  illustrated  by  subjecting  a 
gas  to  varying  pressures,  and  it  will  be  seen  that  when  the  pressure 
is  doubled  the  volume  of  gas  is  reduced  to  one-half,  and  so  on. 

Just  as  in  the  case  of  the  law  of  Charles,  modern  investigations 
have  shown  that  the  law  of  Boyle  is  not  a  mathematical  truth.  It 
is  found  not  to  be  absolutely  true  of  any  gas,  for,  with  the  exception 
of  hydrogen,  all  gases  are  more  compressible  than  is  demanded  by 
the  law.  Hydrogen  deviates  from  the  law  in  an  opposite  sense,  in 
that  it  requires  a  higher  pressure  than  the  law  would  indicate,  in 
order  to  reduce  a  volume  of  it  to  a  given  point.  These  deviations 
from  Boyle's  law  are  explained  by  the  operation  of  two  causes  ; 
first,  the  attraction  exerted  by  gaseous  particles  upon  each  other  ; 
second,  the  fact  that  increased  pressure  diminishes  the  space 
between  the  molecules,  and  not  the  actual  space  occupied  by  the 
molecules  of  a  gas.  When  the  former  cause  predominates,  the 
gas  deviates  from  the  law  by  being  more  compressible  ;  in  the  case 
of  hydrogen  the  second  cause  operates  more  powerfully.  (See 
Kinetic  Theory  of  Gases.)  For  ordinary  purposes  of  calculation 
the  law  of  Boyle  may  be  regarded  as  true. 

As  the  volume  of  a  given  weight  of  gas  is  so  intimately  related 
to  the  pressure,  and  as  the  atmospheric  pressure  is  variable,  it 
becomes  necessary,  in  all  quantitative  manipulation  with  gases,  to 
know  the  actual  pressure  under  which  the  gas  is  at  the  time  of 
measurement,  and  to  refer  the  volume  to  a  standard  pressure. 
The  pressure  that  has  been  adopted  as  the  standard  is  that  of  a 
column  of  mercury  760  mm.  in  height.  (See  Atmosphere.) 

If  v  equals  the  volume  of  gas  measured  at  p  pressure,  and  2/0 
the  volume  at  the  standard  pressure,  then 


In  practice  it  is  most  usual  to  make  both  correction  for  tempe- 


72  Introductory  Outlines 

rature  and  pressure  together  ;  then  ?'0  being  the  volume  at  the 
standard  temperature  and  pressure,  we  get 

^0=i+To~o3665/  '  760  * 

The  Liquefaction  of  Gases.— Under  certain  conditions  of  tem- 
perature and  pressure,  the  law  of  Charles  and  the  law  of  Boyle  both 
completely  break  down.  According  to 
the  law  of  Charles,  100  c.c.  of  a  gas  at 
o°  C.  should  occupy  96.4  c.c.  if  the  tem- 
perature were  lowered  to  —  10°.  If  100 
c.c.  of  the  gas  sulphur  dioxide  at  o°  C. 
be  confined  in  a  glass  tube  standing  in 
mercury,  and  the  gas  be  cooled  to  —  10° 
by  surrounding  the  tube  with  a  freezing 
mixture,  it  will  be  found  that  the  volume 
of  gas,  instead  of  occupying  96.4  c.c., 
has  been  reduced  to  a  few  cubic  centi- 
metres only,  and  that  the  surface  of  the 
mercury  in  the  tube  is  wet  owing  to  the 
presence  of  a  minute  layer  of  a  colourless 
liquid  upon  it.  In  this  case  the  law  of 
Charles  has  broken  down,  and  the  sul- 
phur dioxide  has  passed  from  the  gaseous 
^frV  to  the  liquid  state. 

^_  —^-  r^  Similarly,  according  to  the  law  of 
Boyle,  100  c.c.  of  a  gas  measured  at  the 
standard  pressure  should  occupy  25  c.c. 
when  exposed  to  a  pressure  of  four  additional  atmospheres.  If 
loo  c.c.  of  the  gas  sulphur  dioxide  be  enclosed  in  one  limb  of  a  long 
U-tube,  as  shown  in  Fig.  i,  the  other  limb  being  filled  with  air, 
and  the  two  gases  be  simultaneously  exposed  to  increased  pressure 
by  raising  the  mercury  reservoir,  it  will  be  seen  that  at  first  the 
gases  in  both  tubes  are  compressed  equally.  As  the  pressure 
approaches  three  atmospheres,  however,  the  mercury  will  be  seen 

*  The  student  should  familiarise  himself  with  the  method  of  calculating  the 
changes  of  volume  suffered  by  gases,  by  changes  of  temperature  and  pressure, 
by  working  out  a  number  of  examples  such  as  the  following : — 

1.  If  30  litres  of  gas  are  cooled  from  25°  to  o°,  what  is  the  diminution  in 
volume,  the  pressure  being  constant?     Ans.  2.51  litres. 

2.  If  a  litre  of  air  at  o°  weighs  1.293  grammes  when  the  barometer  is  at 


FIG.  i. 


Liquefaction  of  Gases  73 

to  rise  much  more  rapidly  in  the  tube  containing  the  sulphur 
dioxide,  and  when  the  mercury  reservoir  has  been  raised  to  such  a 
height  that  the  gases  are  subjected  to  four  atmospheres,  the  sulphur 
dioxide  will  have  completely  broken  down,  and  will  be  entirely  con- 
verted into  a  few  drops  of  liquid,  which  appear  upon  the  surface  of 
the  mercury.  The  air  meantime,  in  the  other  limb,  will  be  found  to 
occupy  25  c.c.,  as  that  gas  at  that  pressure  obeys  Boyle's  law  almost 
absolutely.  We  see,  therefore,  that  at  a  certain  temperature  and  at 
a  certain  pressure  the  gas  sulphur  dioxide  begins  rapidly  to  depart 
from  the  laws  of  Charles  and  Boyle,  and  ultimately  passes  into  the 
liquid  condition. 

All  gases,  when  exposed  to  certain  conditions  of  temperature 
and  pressure,  conditions  which  are  special  for  each  different 
gas,  will  pass  from  the  gaseous  to  the  liquid  state  ;  and 
as  the  point  at  which  liquefaction  takes  place  is  approached, 
the  departures  from  Boyle's  law  become  more  and  more  pro- 
nounced. 

The  first  substance,  recognised  as  being  under  ordinary  condi- 
tions a  true  gas,  that  was  transformed  into  the  liquid  condition 
was  chlorine,  which  was  liquefied  in  the  year  1806  by  Northmore. 
The  true  nature  of  this  liquid  was 
not  understood  until  Faraday  inves- 
tigated the  subject. 

In  his  earlier  experiments  Fara- 
day's method  consisted  in  sealing 
into  a  bent  glass  tube  (Fig.  2)  sub- 
stances which,  when  heated,  would 
yield  the  gas  ;  the  substances  being 
contained  in  one  limb  of  the  tube, 
and  the  empty  limb  being  immersed  FIG.  2. 

in  ice.  The  pressure  exerted  by  the  gas  thus  generated  in  a  con- 
fined space  was  sufficient  to  cause  a  portion  of  it  to  condense  to 

760  mm. ,  what  will  be  the  weight  of  a  litre  of  air  at  27°,  the  barometer 
standing  at  the  same  height?  Ans.  1.177  grammes. 

3.  What  will  be  the  weight  of  a  litre  of  air  at  42°  when  the  barometer  stands 
at  735  mm.  ?    Ans.   1.084  grammes. 

4.  Air  at  a  temperature  of  15°  is  enclosed  in  a  vessel  and  heated  to  93°. 
Compare  the  pressure  of  the  enclosed  air  with  that  of  the  atmosphere.     Ans. 
As  61  :  48. 

5.  What  will  be  the  volume,  at  the  standard  temperature  and  pressure,  of 
500  c.c.  of  hydrogen,   measured  at  20°,  and  under  a  pressure  of  800  mm.? 
Arts.  490  c.c. 


74  Introductory  Outlines 

the  liquid  state,  and  the  liquid  collected  in  the  cooled  limb.  In 
this  way  Faraday  liquefied  such  gases  as  chlorine,  sulphur  dioxide, 
ammonia,  cyanogen.  In  his  later  experiments  Faraday  compressed 
the  gas  by  means  of  a  small  compression  pump,  and  at  the  same 
time  applied  a  low  degree  of  cold,  and  by  so  doing  he  succeeded 
in  liquefying  carbon  dioxide,  hydrochloric  acid,  nitrous  oxide,  and 
other  gases.  There  were  a  number  of  gases,  however,  which  Fara- 
day found  it  impossible  to  liquefy,  such  as  hydrogen,  oxygen,  nitro- 
gen, marsh  gas,  nitric  oxide,  carbon  monoxide,  &c.  It  became  the 
custom  to  call  these  permanent  gases >  and  this  term  was  applied  to 
them  until  the  year  1877. 

In  that  year  it  was  proved  by  Pictet,  and  independently  by  Cail- 
letet,  that  under  sufficiently  strong  pressure,  and  a  sufficiently  low 
degree  of  cold,  the  so-called  permanent  gases  could  in  the  same 
way  be  reduced  to  the  liquid  condition.  Pictet's  method  was  in 
principle  the  same  as  that  employed  by  Faraday,  the  difference 
being  that  with  the  machinery  at  his  disposal  he  was  able  to 
employ  enormously  increased  pressure  and  a  greater  degree  of 
cold.  For  the  liquefaction  of  oxygen,  a  quantity  of  potassium 
chlorate  was  heated  in  a  strong  wrought-iron  retort,  to  which  was 
connected  a  long  horizontal  copper  tube  of  great  strength  and  small 
bore.  At  the  extreme  end  of  this  tube  there  was  a  pressure  gauge 
capable  of  indicating  pressures  up  to  800  atmospheres,  and  a  stop- 
cock. The  tube  was  cooled  by  being  contained  in  a  wider  tube, 
through  which  a  constant  stream  of  liquid  carbon  dioxide,  at  a  tem- 
perature of  —  120°  to  —  140°,  was  caused  to  flow. 

The  machinery  employed  to  maintain  this  flow  of  liquefied  car 
bon  dioxide  was  somewhat  elaborate,  consisting  of  condensing  and 
exhaust  pumps  for  liquefying  and  rapidly  evaporating  sulphur 
dioxide,  and  similar  condensing  and  exhaust  pumps  for  liquefying 
and  rapidly  evaporating  carbon  dioxide  :  the  sulphur  dioxide  being 
merely  the  refrigerating  agent  used  to  assist  the  liquefaction  of 
the  carbon  dioxide.  This  machinery  was  driven  by  two  eight- 
horse-power  engines.  As  the  potassium  chlorate  was  heated 
and  oxygen  evolved,  the  internal  pressure  in  the  retort  and 
copper  tube  rapidly  rose,  and  its  amount  was  indicated  by  the 
gauge. 

When  the  stop-cock  upon  the  end  of  the  tube  was  opened,  liquid 
oxygen  was  forcibly  driven  out  in  the  form  of  a  jet. 

In  the  method  employed  by  Cailletet,  the  pressure  to  which  the 
gas  is  subjected  is  obtained  by  purely  mechanical  means.  The 


Liquefaction  of  Gases 


gas  to  be  liquefied  is  introduced  into  a  glass  tube  (Fig.  3),  the 
narrow  end  of  which  consists  of  a  strong  capillary  tube.  The  tube 
carries  a  metal  collar,  which  enables  it  to  be  secured  in  position 
in  the  strong  steel  bottle  (Fig.  4),  by  means  of  a  nut,  E'  (Fig.  5), 
which  screws  into  the  mouth.  The  bottle,  which  is  partially  rilled 
with  mercury,  is  connected  by  means  of  a  flexible  copper  tube  of 
fine  bore  with  a  small  hydraulic  pump,  by  means  of  which  water 
is  forced  into  the  steel  bottle.  The  water  so  driven  in  forces  the 


( 


FIG.  3. 


FIG.  4. 


mercury  up  into  the  glass  tube  T,  and  thereby  compresses  the 
contained  gas.  In  this  way  a  pressure  of  several  hundred  atmos- 
pheres may  be  applied  to  the  gas.  In  his  earlier  experiments 
Cailletet  depended  almost  entirely  for  the  refrigeration  he  required 
upon  the  fact,  that  when  a  gas  is  allowed  suddenly  to  expand  it 
undergoes  a  great  reduction  in  temperature.  This  method  of 
cooling  may  be  termed  internal  refrigeration.  In  the  case  of 
cxygen,  the  gas  was  first  subjected  to  a  pressure  of  300  to  400 


76 


Introductory  Outlines 


B 


atmospheres,  and  was  then  allowed  suddenly  to  expand  by  a  rapid 
release  of  the  pressure.  The  result  of  the  sudden  expansion  was 
to  momentarily  lower  the  temperature  of  the  gas  to  such  a  point 
that  the  tube  was  filled  with  a  fog,  or  mist,  consisting  of  liquid 
particles  of  oxygen. 

This  principle,  namely,  the  self-cooling  of  a  gas  by  its  own 
sudden  expansion,  has  recently  been  applied  for  the  liquefaction 
of  oxygen  in  large  quantities.  When  oxygen  under  considerable 
pressure,  say  120  atmospheres,  is  allowed  to  escape  from  a  fine 
orifice  at  the  end  of  a  long  pipe, -the  issuing  gas  suddenly  expands, 
and  thereby  its  temperature  is  greatly  lowered.  If  this  self-cooled 

gas  is  made  to 
sweep  over  the  pipe 
from  which  it  is 
escaping,  it  will 
cool  the  pipe,  and 
therefore  lower  the 
temperature  of  the 
remaining  gas  be- 
fore it  issues.  In 
this  way  the  cooling 
effect  becomes  cu- 
mulative, the  initial 
temperature  of  the 
gas  before  it  es- 
capes being  con- 
tinually brought 
lower  and  lower, 
until  at  last  the 
point  is  reached  at 
which  the  oxygen 
is  liquefied.* 

If  the  oxygen  be 
first  cooled  to  about 
—  80°  by  means  of 
solid  carbon  di- 
oxide, then  in  a  few 

minutes,  by  the  further  cooling  due  to  its  own  expansion,  the  tem- 
perature will  fall  below  the  boiling-point  of  oxygen,  and  the 
liquefied  gas  be  obtained. 

The  apparatus  for  the  purpose  is  shown  in  Fig.  6.t     Oxygen 

*  Linde,  The  Engineer,  Oct.  4,  1895. 

f  Designed  by  Dewar  ;  constructed  by  Messrs.  Lennox,  Reynolds  &  Fyfe. 


FIG.  6. 


Liquefaction  of  Gases  77 

under  a  pressure  of  120  to  140  atmospheres  is  passed  through  a 
series  of  spirals  of  fine  copper  pipe  contained  in  the  chamber  C, 
which  is  encased  in  a  non-conducting  jacket  of  cork-dust.  The 
gas  enters  by  the  pipe  O  (seen  in  the  enlarged  section),  and  passes 
through  the  spiral  S  S,  which  is  immersed  in  a  mixture  of  alcohol 
and  solid  carbon  dioxide  (the  liquid  carbon  dioxide  from  the  reservoir 
being  admitted  into  the  alcohol  through  the  valve  W,  which  is  regu- 
lated by  the  screw  B).  The  oxygen  thus  cooled  passes  through  the 
double  spiral  pipe  D  D,  which  ultimately  extends  through  the 
bottom  of  the  chamber,  and  terminates  in  a  stirrup,  U,  the  short 
end  of  which  is  closed.  In  the  bend  of  this  stirrup  there  is  a  fine 
hole,  which  can  be  closed  or  opened  at  will  by  the  pointed  end 
of  the  rod  V,  connected  to  the  screw  A.  On  opening  this  valve, 
the  oxygen,  already  cooled  to  about  —  80°,  escapes  from  the  hole 
under  a  pressure  of  120  to  140  atmospheres.  It  instantly  expands, 
and  is  thereby  cooled  still  lower.  This  cold  gas  is  prevented  from 
escaping  at  once  into  the  atmosphere  by  the  glass  tube  G,  but  is 
compelled  to  rush  upwards  (as  shown  by  the  arrows),  and,  sweep- 
ing past  the  double  spiral  D  D,  cools  this  pipe,  and  therefore  the 
succeeding  portions  of  issuing  oxygen.  In  a  few  minutes  the  tem- 
perature of  this  pipe  is  thereby  brought  so  low,  that  the  further 
cooling  of  the  gas  by  its  expansion  causes  the  liquefaction  of  a 
portion  of  it,  and  a  fine  spray  of  liquid  is  seen  to  spurt  out  from 
the  hole.  This  spray  quickly  increases  in  quantity,  and  rapidly 
collects  as  a  clear  liquid  in  the  glass  tube  G.  This  tube  is  double- 
walled,  the  space  between  the  walls  being  perfectly  vacuous.  In 
such  a  vessel  the  liquid  oxygen  may  be  kept  for  a  considerable 
time,  evaporating  only  very  slowly  in  spite  of  its  extremely  low 
boiling-point,  as  it  has  been  found  that  such  a  vacuous  envelope 
forms  the  most  perfect  non-conductor. 

The  instruments  designed  by  Linde  in  Germany,  and  by  Hamp- 
son  in  England,  and  known  as  air-liquefiers,  are  constructed  on 
precisely  similar  principles.  In  this  case,  however,  the  preliminary 
cooling  by  means  of  solid  carbon  dioxide  is  dispensed  with  ;  for 
instead  of  a  limited  and  comparatively  small  supply  of  com- 
pressed gas  in  a  steel  cylinder,  an  unlimited  supply  of  air  is 
delivered  into  the  machine,  under  a  pressure  of  120  to  160 
atmospheres,  by  means  of  powerful  compression  pumps  driven 
by  an  engine. 

By  an  extension  of  the  same  principles  hydrogen  was  first  suc- 
cessfully liquefied  in  1898.  In  this  case,  however,  the  gas  requires 


Introductory  Outlines 


to  be  previously  cooled  to  about  —  200°  before  expansion  is  allowed  to 
take  place.  By  utilising  the  low  temperatures  which  can  be  obtained 
by  means  of  boiling  liquefied  gases,  it  has  now  become  possible 
to  liquefy  all  the  known  gases  by  cold  alone,  that  is,  without  the 
application  of  pressure  ;  in  other  words,  their  temperatures  can 
be  brought  down  below  their  boiling-points,  under  which  circum- 
stances they  must  obviously  assume  the  liquid  state.  For  example, 
liquefied  ethylene  boils  at  -  103.5°  J  if»  therefore,  a  stream  of  nitrous 
oxide  is  passed  through  a  tube  immersed  in  a  bath  of  liquid  eihy- 
lene,the  nitrous  oxide  will  be  cooled  below  its  boiling-point  (  —  89.8°)., 
and  will  consequently  be  reduced  at  once  to  the  liquid  state. 
Again,  liquid  oxygen  boils  at  —182.5°. 
This  boiling  liquid  therefore  is  sufficiently 
cold  to  cool  marsh  gas  below  its  boiling- 
point,  namely,  — 164.7°,  and  therefore  to 
cause  its  liquefaction. 

^,v  Moreover,  by  the  rapid  evaporation  of  liquid 
oxygen  the  temperature  may  readily  be 
lowered  to  the  point  at  which  air  will  liquefy. 
.-•°  Thus,  if  a  quantity  of  liquid  oxygen  in  the 
glass  tube  O  (Fig.  7),  which  is  provided  with 
a  vacuous  envelope,  V,  be  made  to  boil 
,-n  rapidly  by  putting  the  pipe  P  in  connection 
with  an  exhaust  -  pump,  the  temperature 
quickly  falls  to  —  200°,  when  air  itself  be- 
comes liquefied  without  the  application  of 
pressure  ;  and  drops  of  liquid  air  quickly 
collect  upon  the  walls  of  the  inner  empty  tube, 
N,  which  is  freely  open  to  the  atmosphere. 
In  this  way  considerable  quantities  of  lique- 
fied air  can  be  collected  in  a  few  minutes. 
By  means  of  boiling  liquid  hydrogen  the  low  temperature  of 
—  253°  has  been  reached,  at  which  temperature  all  other  known 
gases,  except  helium,  are  frozen  to  the  solid  state.  The  lowest 
temperature  yet  obtained  by  the  rapid  evaporation  of  solid 
hydrogen  is  -260°  (Dewar). 

The  Critical  Point— As  far  back  as  the  year  1869,  it  was 
shown  by  Andrews  that  when  liquid  carbon  dioxide  was  heated 
to  a  particular  temperature,  it  passed  from  the  liquid  to  the  gaseous 
state,  and  that  no  additional  pressure  was  able  to  condense  it  again 
so  long  as  the  temperature  remained  at  or  above  that  point.  This 


FIG.  7. 


Critical  Temperature   of  Gases  79 

particular  temperature  is  called  the  critical  point,  or  the  critical 
temperature  of  the  gas.  In  the  case  of  carbon  dioxide  this  critical 
temperature  is  31.35°,  and  in  order  that  this  gas  may  be  liquefied  by 
pressure,  it  is  an  essential  condition  that  the  temperature  be  below 
that  point  ;  above  32°  no  pressure  is  capable  of  bringing  about 
liquefaction.  All  gases  have  a  critical  temperature,  which  is  special 
for  each  gas,  and  until  the  temperature  of  the  gas  be  lowered  to 
that  point,  liquefaction  is  impossible.  The  critical  temperatures 
of  the  different  gases  vary  through  a  very  wide  range  :  thus, 
the  critical  temperature  of  hydrogen  is  as  low  as  —238°,  while  that 
of  sulphur  dioxide  is  +155.4°.  In  the  third  column  of  the  table 
of  physical  constants  on  page  80  the  critical  temperatures  of  a 
number  of  the  more  common  g~ases  are  given.* 

The  gases  in  this  list,  from  ethylene  downwards,  all  have  their 
critical  temperatures  so  high  that  there  is  no  difficulty  in  cooling 
them  below  these  points.  These  are  the  gases  which  were  first 
reduced  to  the  liquid  state.  The  first  five  upon  the  list  have 
very  low  critical  temperatures  ;  these  are  the  very  gases  which  for 
so  long  resisted  all  attempts  to  liquefy  them,  and  which  were  on 
that  account  called  permanent  gases.  We  now  know  that  the 
failure  to  obtain  them  in  the  liquid  state  was  owing  to  the  fact 
that  the  relation  between  the  critical  temperature  and  the  point 
of  liquefaction  was  not  fully  'ealised.  Just  as  carbon  dioxide 
cannot  be  liquefied  unless  its  temperature  be  brought  down  to 
31.35°,  so  oxygen  resists  liquefaction  under  the  highest  possible 
pressures,  until  its  temperature  be  lowered  to  — 118.8°,  the  critical 
temperature  of  oxygen. 

The  critical  temperature  of  a  gas  is  sometimes  spoken  of  as  the 
absolute  boiling-point. 

Critical  Pressure.— The  particular  pressure  that  is  required 
to  liquefy  a  gas  at  its  critical  temperature  is  called  the  critical 
pressure.  Thus  the  pressure  necessary  to  liquefy  oxygen,  when 
the  temperature  has  been  lowered  to  — 118.8°,  is  58  atmospheres  ; 
while  that  required  to  condense  chlorine  at  its  critical  point,  viz., 
+  141°,  is  84  atmospheres.  At  temperatures  below  the  critical 
temperatures  a  gas  liquefies  under  less  pressure  than  the  critical 

*  For  the  constants  for  the  gases  of  the  Argon  family  see  page  271.  It  may 
be  well  to  remind  the  student  that  such  constants  as  are  here  tabulated 
are  obtained,  from  measurements  involving  very  great  experimental  difficulties, 
and  that  consequently  they  are  always  liable  to  revision.  The  values  here 
given  are  from  the  most  recent  determinations. 


8o 


Introductory  Outlines 


pressure,  until  when  the  temperature  is  lowered  to  the  boiling- 
point  of  the  gas  it  passes  into  the  liquid  state  without  the  applica- 
tion of  any  external  pressure.  The  following  table  contains  the 
most  recently  determined  physical  constants  of  a  number  of 
common  gases  : — 

TABLE  OF  PHYSICAL  CONSTANTS. 


Boiling- 
Pol  nt. 

Melting- 
Point. 

Critical 
Temp. 

Critical 
Pressure. 

Density 
at  Boiling- 
Point. 

Hydrogen 

-253° 

-257° 

-238° 

15-3'Ats. 

0.06 

Nitrogen  . 

-19B-S° 

-213° 

-149° 

27-5 

0.791 

Carbon  monoxide 

-190° 

-207° 

-136° 

33-5 

Oxygen 

-  182.5° 

-223° 

-118.8° 

58.0 

1.131 

Methane  (marsh  gas) 

-164.7° 

-184° 

-   82° 

55-8 

0.416 

Ethylene   . 

-  1°3-  5° 

-169° 

+     9° 

58.0 

0.571 

Nitrous  oxide    . 

-   89-8° 

—  102.7° 

+   37° 

Acetylene  . 

-   827° 

+  35° 

61 

Carbon  dioxide 

-  80° 

+  31-35° 

72.3 

Ammonia  . 

-   38-5° 

-   75-5° 

+  131° 

H3 

Chlorine    . 

-   33-4° 

-}•  141 

84 

i-5°7 

Sulphur  dioxide 

-     10° 

-<-i55-4° 

78.9 

From  the  figures  given  in  this  table  it  will  be  seen  that  the 
critical  pressure  (which  is  the  pressure  required  to  liquefy  a  gas  at 
the  highest  temperature  at  which  pressure  can  possibly  cause  lique- 
faction) is  in  most  cases  comparatively  small.  In  only  one  instance, 
namely,  ammonia,  is  it  over  loo  atmospheres,  and  falling  in  the 
case  of  hydrogen  as  low  as  15.3  atmospheres.  The  enormous 
pressures,  therefore,  amounting  often  to  many  hundred  atmospheres, 
which  some  of  the  earlier  experimenters  employed  in  attempting  to 
effect  the  liquefaction  of  the  so-called  permanent  gases,  are  thus 
seen  to  have  been  efforts  in  an  entirely  wrong  direction.  It  was  not 
greater  pressure  that  was  required,  but  the  means  of  cooling  the 
gases  to  a  sufficiently  low  temperature. 

In  ordinary  language  such  a  gas  as  chlorine  is  spoken  of  as  ai\ 
easily  liquefied  gas,  while  oxygen  would  be  described  as  a  difficultly 
liquefied  gas.  Strictly  speaking,  however,  and  considering  them 
from  a  comparable  standpoint,  it  would  perhaps  be  more  correct 
to  regard  them  in  exactly  the  opposite  light.  Thus,  taken  at  their 
respective  critical  temperatures,  oxygen  is  liquefied  by  a  pressure 
of  58  atmospheres  ;  while  at  the  critical  temperature  of  chlorine  this 


Diffusion  of  Gases  81 

gas  requires  a  pressure  of  84  atmospheres  to  reduce  it  to  the  liquid 
state.  At  o°  it  is  true  chlorine  may  be  liquefied  by  a  pressure  of 
only  6.  atmospheres,  but  it  must  be  remembered  that  o°  is  141 
degrees  below  the  critical  temperature  of  this  gas.  Long  before 
oxygen  has  been  cooled  141  degrees  below  its  critical  temperature, 
which  would  be  down  to  —254°,  it  not  only  passes  into  the  liquid 
state  without  the  application  of  any  external  pressure  at  all,  but  is 
frozen  to  the  solid  state. 

Diffusion  of  Gases.— If  a  jar  filled  with  hydrogen  be  placed 
mouth  to  mouth  with  a  jar  of  air,  the  hydrogen  being  uppermost, 
it  will  be  found  that  after  the  lapse  of  a  few  minutes  some  of  the 
hydrogen  will  have  passed  into  the  bottom  jar  containing  air,  and 
some  of  the  air  will  have  made  its  way  up  into  the  hydrogen  jar. 
The  light  gas  hydrogen  does  not,  as  might  have  been  supposed, 
remain  floating  upon  the  air,  which  is  14.44  times  as  heavy,  but 
gradually  escapes  into  the  lower  jar  ;  and  the  heavier  gas  finds  its 
way,  in  opposition  to  gravitation,  into  the  upper  jar.  This  process 
goes  on  until  there  is  a  uniform  mixture  of  air  and  hydrogen  in  both 
jars,  and  the  gases  never  separate  again  according  to  their 
densities. 

This  transmigration  of  gases  will  take  place  even  through  tubes 
of  considerable  length :  thus,  if  two  soda-water  bottles  be  filled  one 
with  hydrogen  and  the  other  with  oxygen,  and  the  two  bottles  be 
connected  by  a  piece  of  glass  tube  a  metre  in  length,  the  system 
being  held  in  a  vertical  position  with  the  light  hydrogen  upper- 
most, it  will  be  found  after  an  hour  or  two  that  the  two  gases 
have  become  mixed.  Some  of  the  hydrogen  will  have  descended 
through  the  long  tube  into  the  lower  bottle,  and  in  like  manner 
a  portion  of  the  oxygen,  although  nearly  sixteen  times  as  heavy 
as  hydrogen,  will  have  travelled  up  into  the  top  bottle.  That  the 
gases  have  so  mixed  may  be  readily  shown  by  applying  a  lighted 
taper  to  the  mouth  of  each  bottle,  the  detonation  which  then  takes 
place  proving  that  the  bottles  contain  a  mixture  of  oxygen  and 
hydrogen.  This  passage  of  one  gas  into  another  is  called  the 
diffusion  of  gases.  It  was  observed  by  Graham  that  when  the 
two  gases  were  separated  from  each  other  by  a  thin  porous 
septum,  such,  for  instance,  as  a  piece  of  unglazed  porcelain  (so- 
called  "  biscuit "),  or  plaster  of  Paris,  the  pressure  of  the  gas  on 
the  two  sides  of  the  porous  partition  did  not  remain  the  same 
during  the  process  of  diffusion  :  that  is  to  say,  one  gas  made  its 
way  through  the  partition  faster  than  the  other,  and  it  was  noticed 

F 


82 


Introductory  Outlines 


that  the  lighter  the  gas  the  more  rapidly  was  it  able  to  transpire 
or  diffuse  through  the  porous  medium.  This  fact,  viz.,  that  a  light 
gas  diffuses  more  rapidly  than  a  heavier  one,  may  be  observed 
in  a  variety  of  ways.*  The  apparatus  seen  in  Fig.  8  is  a  modified 
form  of  Graham's  diffusiometer.  It  consists  of  a  long  glass  tube 
with  an  enlargement  or  bulb  near  to  one  end.  Into  the  short  neck 
of  this  bulb  there  is  fastened  a  thin  diaphragm  of  stucco,  or  other 
porous  material.  If  the  apparatus  be  filled  with  hydrogen  by  dis- 
placement, the  short  neck  being  closed  by  a  cork,  and  the  long 
limb  be  immersed  in  water,  it  will  be  seen,  upon  the  withdrawal 


FIG.  8. 


FIG.  9. 


of  the  cork,  that  the  water  rapidly  rises  in  the  long  tube.  The 
hydrogen  diffusing  out  through  the  diaphragm  so  much  more 
rapidly  than  air  can  make  its  way  in,  a  diminution  in  pressure 
within  the  apparatus  results,  and  this  causes  the  water  to  ascend 
in  the  tube.  The  same  phenomenon  may  be  seen  even  more 
strikingly  by  means  of  the  apparatus,  Fig.  9,  which  consists  of  a 
tall  glass  U-tube,  upon  the  end  of  one  limb  of  which  there  is 
fastened,  by  means  of  a  cork,  a  porous  cylindrical  pot,  such  as 

*  See  Experiments  Nos.  350-359, Newth's  "Chemical  Lecture  Experiments," 
new  ed. 


Diffusion  of  Gases  83 

is  used  in  an  ordinary  Bunsen  battery.  The  U-tube  is  half 
filled  with  coloured  water.  Under  ordinary  circumstances  air  is 
continually  diffusing  through  the  porous  pot,  but  as  it  passes  at 
an  equal  rate  in  both  directions,  there  is  no  disturbance  of  the 
pressure,  and  consequently  the  coloured  water  remains  level  in 
the  two  limbs.  If  now  a  beaker  containing  hydrogen  be  brought 
over  the  apparatus,  as  seen  in  the  figure,  the  hydrogen  will  stream 
through  the  porous  pot  so  much  more  rapidly  than  the  air  in  the 
pot  can  make  its  way  out,  that  there  will  be  an  increase  in  the 
total  amount  of  gas  inside  the  apparatus,  which  will  be  instantly 
rendered  evident  by  the  change  of  level  of  the  liquid  in  the  U-tube, 
the  water  being  forcibly  driven  down  the  tube  which  carries  the 
porous  pot.  Upon  removing  the  beaker  the  reverse  operation 
will  at  once  take  place  ;  the  hydrogen  inside  the  apparatus  now 
rapidly  diffuses  out,  and  much  more  quickly  than  air  can  pass  in, 
consequently  a  reduction  of  pressure  within  the  apparatus  results, 
which  is  indicated  by  a  disturbance  of  the  level  of  the  water 
in  the  tube,  in  the  opposite  direction  to  that  which  occurred  at 
first. 

The  Law  Of  Gaseous  Diffusion.— Graham  established  the  law 
according  to  which  the  diffusion  of  gases  is  regulated,  and  it  may 
be  thus  stated  :  The  relative  velocities  of  diffusion  of  any  two 
gases  are  inversely  as  the  square  roots  of  their  densities. 

The  density  of  hydrogen  being  I,  that  of  air  is  14.44,  the  velocity 
of  the  diffusion  of  hydrogen,  therefore,  as  compared  with  that  of 
air,  will  be  in  the  ratio  of  */M44  to  v/i.  VM44  =  3-8,  \/i  =  i. 
Therefore  hydrogen  diffuses  3.8  times  faster  than  air  ;  or  3.8  volumes 
of  hydrogen  will  pass  out  through  a  porous  septum,  while  only  I 
volume  of  air  can  enter. 

If  d  =  the  density  of  a  gas,  air  being  unity,  and  i>  =  the  volume 
of  the  gas  which  diffuses  in  the  same  time  as  I  volume  of  air,  then 


The  following  table  gives  in  the  last  column  the  results  obtained 
by  Graham,  which  will  be  seen  to  accord  very  closely  with  the  cal- 
culated numbers  demanded  by  the  law  of  diffusion  : — 


Introductory  Outlines 


, 

Volume  of  Gas 

Density  of  Gas 

/- 

which  Diffused  in 

Name  of  Gas. 

compared  with 

-V  - 

the  same  Time  as 

A\r  —  d. 

d 

one  Volume  of 

Air. 

Hydrogen 

o.  06926 

3-7794 

3-83 

Marsh  gas 

°-559 

1-3375 

1-344 

Carbon  monoxide      . 

0.9678 

1.0165 

1.0149 

Nitrogen  . 

0.9713 

1.0147 

1.0143 

Oxygen      . 
Sulphuretted  hydrogen 

1.1056 
1.1912 

0.9510 
0.9162 

0.9487 
0.95 

Carbon  dioxide 

1.5290 

o.  8087 

0.812 

Sulphur  dioxide 

2.247 

0.6671 

0.68 

The  property  of  diffusion  is  sometimes  made  use  of  in  order  to 
separate  gases,  having  different  densities,  from  gaseous  mixtures. 
This  process  of  separation  by  diffusion  is  known  as  atmolysis. 
The  principle  may  readily  be  illustrated  by  causing  a  mixture  of 

oxygen  and  hydrogen,  in 
proportion  to  form  an  ex- 
plosive mixture,  to  slowly 
traverse  tubes  made  of 
porous  material,  such  as 
ordinary  tobacco  pipes. 
Two  such  pipes  may  be 
arranged  as  shown  in  Fig. 
10,  and  the  gaseous  mix- 
ture passed  through  in  the 
direction  indicated  by  the 
arrow.  On  collecting  the 

*  .  . 

issuing  gas  over  water  in  a 
pneumatic  trough,  it  will 
be  found  to  have  so  far 
lost  the  hydrogen,  by  dif- 

FlG  la  fusion  through    the    tube, 

that  a   glowing  splint    of 
wood  when  introduced  into  it  will  be  reignited. 

From  the  rate  of  diffusion  of  ozone,  in  a  mixture  of  ozone  and 
oxygen,  Soret  was  able  to  calculate  the  density  of  this  allotropic 
form  of  oxygen,  and  so  confirm  the  result  he  had  previously  ob- 
tained by  other  methods  (see  Ozone). 

Attempts  have  been  made  to  utilise  this  principle  in  order  to 
obtain  oxygen  from  the  air.  The  relative  densities  of  oxygen  and 


The  Kinetic  Theory  85 

nitrogen  are  as  1 6  to  14  ;  the  rate  of  diffusion,  therefore,  of  nitrogen 
is  slightly  greater  than  that  of  oxygen. 

Effusion  is  the  term  applied  by  Graham  to  the  passage  of  gases 
through  a  fine  opening  in  a  very  thin  wall,  and  he  found  that  it 
followed  the  same  law  as  diffusion.  Bunsen  utilised  this  principle 
for  determining  the  density,  and  therefore  the  molecular  weights, 
of  certain  gases.  The  method,  in  essence,  is  as  follows  : — A 
straight  glass  eudiometer  is  so  constructed,  that  a  gas  contained 
in  it  can  be  put  into  communication  with  the  outer  air  through  a 
minute  pin-hole  in  a  thin  platinum  plate.  The  gas  is  confined  in 
the  tube,  which  is  placed  in  a  cylindrical  mercury  trough,  by 
means  of  a  stop-cock  at  the  top.  When  the  tube  is  depressed 
in  the  mercury,  and  the  cock  opened,  the  gas  escapes  through 
the  minute  perforation  in  the  platinum  plate,  and  its  rate  of  effu- 
sion is  determined  by  the  time  occupied  by  a  glass  float  placed 
in  the  tube  in  rising  a  graduated  distance  within  the  eudiometer. 

The  flow  of  gases  through  capillary  tubes  is  called  transpiration 
of  gases.  In  this  case  the  friction  between  the  gas  and  the  tubes 
becomes  a  factor  in  the  movement,  so  that  this  phenomenon  is 
not  governed  by  the  same  law  as  gaseous  diffusion. 

The  Kinetic  Theory  of  Gases.— The  term  kinetic  signifies 
motion,  and  as  applied  to  this  theory  it  expresses  the  modern 
views  of  physicists  concerning  matter  in  the  gaseous  state,  and 
serves  to  harmonise  and  explain  the  physical  laws  relating  to 
the  properties  of  gases.  Matter  in  the  state  of  gas  or  vapour 
is  regarded  as  an  aggregation  of  molecules  in  which  the  attractive 
forces  which  tend  to  hold  them  together  are  reduced  to  a  minimum, 
and  in  which  the  spaces  that  separate  them  are  at  a  maximum. 
These  molecules  are  in  a  state  of  rapid  motion,  each  one  moving 
in  a  straight  line  until  it  strikes  some  other  molecule,  or  rebounds 
from  the  walls  of  the  containing  vessel,  when  it  continues  its  move- 
ment in  another  direction  until  it  is  once  more  diverted  by  another 
encounter.  As  they  constantly  encounter  and  rebound  from  each 
other,  it  will  be  evident  that  at  any  given  instant  some  will  be 
moving  with  a  greater  speed  than  others  ;  the  majority,  however, 
will  have  an  average  velocity.  In  these  encounters  no  loss  of 
energy  results  so  long  as  the  temperature  remains  constant,  but 
any  change  of  temperature  results  in  a  change  in  the  velocity  of 
movement  of  the  molecules,  the  speed  being  increased  with 
increased  heat.  The  actual  volume  of  the  molecules  is  very  small 
as  compared  with  the  space  occupied  by  the  mass  ;  the  space 


86  Introductory  Outlines 

between  the  molecules,  therefore,  in  which  they  pass  to  and  fro, 
is  relatively  very  great.  As  the  molecules  are  constantly  colliding 
and  rebounding,  the  distances  between  them,  as  well  as  their  speed, 
will  be  sometimes  greater  and  sometimes  less  ;  but  there  will  be 
an  average  distance,  which  is  known  as  the  mean  free  path  of  the 
molecule. 

The  pressure  exerted  by  a  gas,  or  its  elastic  force,  is  the  combined 
effect  of  the  bombardment  of  its  molecules  against  the  containing 
vessel  ;  in  other  words,  the  pressure  of  a  gas  is  proportional  to  the 
sum  of  the  products  obtained  by  multiplying  the  mass  of  each 
molecule  by  half  the  square  of  its  velocity.  It  will  be  obvious 
that  if  the  space  within  which  a  given  mass  of  gas  is  confined  be 
reduced,  the  number  of  impacts  of  the  molecules  against  the  walls 
of  the  containing  vessel,  in  a  given  time,  will  be  increased,  and 
therefore  the  pressure  it  exerts,  or  its  elastic  force,  will  also  be 
increased.  If  the  space  be  reduced  to  one-half  the  original,  the 
number  of  these  impacts  will  be  doubled,  or  in  other  words,  the 
number  of  impacts  in  a  given  time  is  inversely  as  the  volume. 
This  statement  is  simply  the  law  of  Boyle  stated  in  the  language 
of  the  kinetic  theory. 

When  a  given  mass  of  gas  contained  in  a  confined  space  is 
heated,  the  pressure  it  exerts,  or  its  elastic  force,  is  increased.  But 
as  the  number  of  molecules  present  has  not  been  increased  by 
raising  the  temperature  of  the  gas  (provided  no  chemical  decom- 
position of  the  gas  is  brought  about  by  the  change  of  temperature), 
the  increased  pressure  can  only  have  resulted  from  the  greater 
frequency,  and  greater  energy,  of  the  impacts  of  the  molecules 
against  the  walls  of  the  vessel,  owing  to  their  greater  velocity. 

Two  equal  volumes  of  different  gases  under  the  same  conditions 
of  temperature  and  pressure,  exert  the  same  elastic  force  upon  the 
containing  vessels,  that  is  to  say,  the  kinetic  energy  in  each  volume 
is  the  same.  According  to  Avogadro's  hypothesis,  equal  volumes 
of  all  gases  under  the  same  conditions  of  temperature  and  pressure 
contain  an  equal  number  of  molecules,  however  much  the  weight 
of  these  molecules  may  vary  ;  therefore  the  average  kinetic  energy 
of  each  individual  molecule  will  be  the  same.  It  follows  from  this 
that  the  mean  velocities  of  different  molecules  must  vary,  and  the 
calculated  numbers  representing  the  actual  velocities  of  movement 
of  the  molecules  of  different  gases  show  that  these  rates  are  pro- 
portional to  the  inverse  square  roots  of  their  respective  densities. 
But  according  to  the  law  of  gaseous  diffusion  (Graham's  law),  the 


The  Kinetic  Theory  87 

relative  rapidity  of  diffusion  of  gases  is  inversely  proportional  to 
the  square  roots  of  their  densities,  hence  by  purely  mathematical 
processes,  based  upon  the  kinetic  theory  of  gases,  the  law  of 
gaseous  diffusion  is  proved  to  be  true.  Similarly,  the  kinetic  theory 
is  applicable  to  the  consideration  of  the  phenomena  of  evaporation 
and  condensation  (see  page  126),  and  to  the  processes  of  solution 
(page  148). 

The  deviations  from  the  laws  of  Boyle  and  Charles,  already 
referred  to,*  are  also  explained  by  the  dynamical  theory  of  gases, 
from  considerations  of  the  following  order  : — 

1.  That  the  molecules  themselves  are  not  mathematical  points, 
but  occupy  a  space  ;  in  other  words,  the  space  occupied  by  the 
actual  particles  of  matter  is  not  infinitely  small  as  compared  with 
the  entire  volume  of  the  gas,  i.e.  the  bulk  of  the  particle  plus  the 
intermolecular  spaces. 

While  the  pressure  upon  a  gas  is  only  slight,  and  therefore  the 
total  volume  occupied  by  a  given  mass  of  the  gas  is  great,  the  bulk 
of  the  actual  particles  themselves  becomes  a  vanishing  quantity  in 
comparison  with  the  total  volume  (i.e.  the  space  occupied  by 
particles,  plus  the  intermolecular  spaces),  and  the  gas  under  these 
circumstances  tends  to  approach  more  nearly  to  the  conditions  of 
an  ideal  gas.  But  when  the  pressure  is  increased,  and  the  total 
volume  thereby  greatly  reduced,  then  the  bulk  of  the  particles 
themselves  begins  to  bear  an  appreciable  proportion  to  the  total 
volume  occupied  by  the  gas. 

2.  That  the  impact  of  the  molecules  against  each  other  and 
against  the  containing  envelope  occupies  time  ;  or,  in  other  words, 
the  time  occupied  by  the  impacts  is  not  infinitely  small  compared 
with  the  time  elapsing  between  the  impacts. 

3.  That  the  molecules  themselves  are  not  entirely  without  attrac- 
tion for  each  other  ;  that  is  to  say,  although  the  attractive  force 
between  the  molecules  which  holds  them  together  in  the  liquid 
and  solid  states  of  matter  is  at  a  minimum  in  the  case  of  gases,  it 
is  not  entirely  absent. 

*  See  page  71. 


CHAPTER   X 

DISSOCIATION— REVERSIBLE   OR  BALANCED 
ACTIONS 

DISSOCIATION  is  the  term  employed  to  denote  a  special  class  of 
chemical  decompositions.  When  potassium  chlbrate  is  heated  it 
breaks  up  into  potassium  chloride  and  oxygen,  thus — 

2KC1O3  =  2KC1  +  302, 

and  when  calcium  carbonate  (chalk)  is  heated  it  breaks  up  into 
calcium  oxide  (lime)  and  carbon  dioxide— 

CaCO3  =  CaO  +  CO2. 

In  the  first  case  the  oxygen  is  incapable  of  reuniting  with  the 
potassium  chloride,  but  in  the  second,  the  carbon  dioxide  can 
recombine  with  the  lime  and  reproduce  calcium  carbonate  :  there- 
fore both  the  following  expressions  are  possible — 

CaCO3  =  CaO  +  CO2, 
and 

CaO  +  CO2  =  CaCO3. 

Reactions  of  this  order  are  known  as  reversible  or  balanced  actions, 
and  the  breaking  up  of  calcium  carbonate  by  the  action  of  heat  is 
termed  dissociation,  while  that  of  the  potassium  chloride  under 
similar  circumstances  is  simple  decomposition. 

When  ammonia  is  passed  through  a  tube  heated  to  a  dull  red 
heat,  the  gas  is  decomposed  into  nitrogen  and  hydrogen — 

2NH3  =  N2  +  3H2, 

and  the  two  gases  pass  out  of  the  heated  tube  as  separated  gases, 
and  do  not  recombine  again.* 

But  when  steam  is  strongly  heated  it  is  dissociated  into  oxygen 

*  Nitrogen  and  hydrogen  can  be  caused  to  unite  under  suitable  conditions 
(see  Ammonia). 


Dissociation  89 

and  hydrogen,  and  as  these  separated  gases  pass  away  from  the 
heated  region  they  reunite,  forming  molecules  of  water  vapour. 
Such  a  reversible  reaction  may  be  thus  expressed — 

2H2O  ;±  2H2  +  O2. 

Again,  when  the  gases  ammonia  and  hydrochloric  acid  are  brought 
together  at  the  ordinary  temperature,  they  unite  to  form  solid 
ammonium  chloride,  and  when  ammonium  chloride  is  heated  it 
dissociates  into  its  two  generators,*  hence  we  have  the  expression — 

NH3  +  HC1  Z.  NH4C1. 

The  corresponding  compound  containing  phosphorus  in  the  place 
of  nitrogen  dissociates  at  a  temperature  as  low  as  —20°,  hence 
when  hydrogen  phosphide  and  hydrochloric  acid  are  mixed  at 
ordinary  temperatures  no  combination  takes  place,  the  separate 
molecules  are  in  the  same  relation  to  one  another  as  those  of 
ammonia  and  hydrochloric  acid  at  a  high  temperature.  When, 
however,  the  mixture  of  gases  is  cooled  below  —  20°,  union  takes 
place  and  crystals  of  phosphonium  chloride  are  formed,  which  at 
once  begin  to  dissociate  into  the  original  gases  as  the  temperature 
again  rises.  The  change,  as  before,  may  be  represented  as  a 
reversible  one — 

PH3  +  HC1  £  PH4CL 

In  such  cases  of  dissociation  as  that  of  calcium  carbonate,  where 
one  of  the  products  is  gaseous  and  the  other  solid,  no  difficulty 
exists  in  separating  the  simpler  compounds  that  result  from  the 
decomposition  ;  but  where  the  products  are  entirely  gaseous,  special 
methods  have  to  be  adopted  to  withdraw  the  one  from  the  other, 
while  they  still  exist  as  separate  molecules,  and  before  they  reunite 
again.  One  such  method,  which  is  well  adapted  for  the  quali- 
tative illustration  of  dissociation,  is  based  on  the  law  of  gaseous 
diffusion.  If  when  ammonium  chloride  is  heated  it  is  dissociated 
into  ammonia,  NH3,  and  hydrochloric  acid,  HC1,  these  two  gases, 
having  the  relative  densities  of  8.5  and  18.25,  will  diffuse  through 
a  porous  medium  at  very  different  rates.  According  to  the  law  of 
diffusion,  these  rates  will  be  inversely  as  the  square  roots  of  the 
densities  of  the  gases  ;  if,  therefore,  the  conditions  are  so  arranged 

*  Baker  has  shown  (May  1894)  that  when  absolutely  dry,  these  gases  do  not 
combine  ;  and  also,  that  when  aqueous  vapour  is  entirely  absent,  ammonium 
chloride  does  not  undergo  this  dissociation. 


9C  Introductory  Outlines 

that  the  heating  of  the  ammonium  chloride  takes  place  in  the 
neighbourhood  of  a  porous  diaphragm,  more  of  the  light  ammonia 
gas  will  diffuse  through  in  a  given  time  than  of  the  heavier  hydro- 
chloric acid,  so  that  a  partial  separation  of  these  gases  will  be 
effected.  Fig.  1 1  shows  a  convenient  arrangement  for  carrying  out 
the  experiment.  A  fragment  of  ammonium  chloride  is  heated  in  a 
short  glass  tube  through  which  passes  the  stem  of  an  ordinary  clay 
tobacco  pipe.  As  the  dissociation  takes  place,  both  of  the  gaseous 
products  begin  to  diffuse  into  the  interior  of  the  porous  clay  pipe, 
but  owing  to  their  greater  rate  of  diffusion,  a  larger  number  of  am- 
monia molecules  will  pass  in,  than  of  hydrochloric  acid,  in  the  same 
time  ;  consequently,  when  the  gases  pass  away  from  the  heated 
region  and  once  more  recombine,  there  will  be  a  surplus  of  am- 
monia molecules  within  the  porous  pipe,  and  for  the  same  reason 
an  excess  of  hydrochloric  acid  molecules  outside.  If  the  gaseous 
contents  of  the  porous  tube  be  driven  out  by  means  of  a  stream  of 


FIG.  ii. 

air  from  an  ordinary  bellows,  the  presence  of  the  free  ammonia  may 
be  recognised  by  allowing  the  air  to  impinge  upon  a  piece  of  paper, 
coloured  yellow  with  turmeric,  which  is  instantly  turned  brown  by 
ammonia.  The  excess  of  hydrochloric  acid  within  the  glass  tube 
may  also  be  proved  by  placing  a  piece  of  blue  litmus  paper  in  the 
tube  before  heating  the  compound,  and  it  will  be  reddened  by  the 
free  hydrochloric  acid. 

In  all  cases  of  dissociation  we  may  imagine  two  opposing  forces 
in  operation,  one  being  the  external  force  supplying  the  energy 
which  tends  to  bring  about  the  disruption  of  the  molecules,  and 
the  other  being  the  force  of  the  chemical  affinity  existing  between 
the  disunited  portions  of  the  molecule,  which  tends  to  bring  about 
their  reunion.  When  these  forces  are  equally  balanced,  the  same 
number  of  molecules  are  dissociated  as  are  recombined  in  a  given 


Dissociation  9 1 

unit  of  time,  and  the  system  is  said  to  be  in  a  state  of  equilibrium. 
If  by  any  means  the  balance  between  the  two  opposing  forces  is 
disturbed,  by  augmenting  or  lessening  either  one  or  the  other  of 
them,  the  equilibrium  of  the  system  will  also  be  disturbed  and  a 
new  condition  of  equilibrium  will  be  set  up,  in  which  again  an  equal 
number  of  molecules  undergo  dissociation  and  combination  in  a 
given  time,  but  in  which  the  ratio  of  the  number  of  united  and  dis- 
united molecules  is  different  from  that  which  obtained  under  the 
former  condition  of  equilibrium.  The  relation  between  these  two 
forces  may  be  most  readily  disturbed,  by  either  a  change  of  tempe- 
rature or  pressure.  Thus,  in  the  case  of  nitrogen  peroxide,  N2O4, 
when  this  gas  is  at  a  temperature  of  26.7°,  20  per  cent,  of  it  is 
dissociated  into  molecules  having  the  composition  NO2  ;  and  so 
long  as  this  temperature  is  maintained  this  ratio  of  the  weight  of 
the  dissociated  molecules  to  the  total  weight  of  the  system  (known 
as  the  fraction  of  dissociation}  still  subsists. 

When  the  temperature  of  the  gas  is  raised  to  60.2°,  the  state  of 
equilibrium  existing  at  the  lower  temperature  is  disturbed,  and  the 
system  gradually  assumes  a  new  condition  of  equilibrium,  where 
once  more  the  actual  number  of  molecules  undergoing  dissociation 
and  recombination  in  a  given  unit  of  time  is  the  same,  but  where 
the  percentage  of  dissociated  molecules  in  the  gaseous  mixture  is 
now  52.04. 

It  might  at  first  be  supposed  when  such  a  gas  is  heated,  and  a 
temperature  is  reached  at  which  the  molecules  are  dissociated,  that 
they  would  all  dissociate,  and  that  the  process  once  begun  would 
rapidly  proceed  until  the  decomposition  was  complete  ;  instead  of 
which,  we  find  a  definite  fraction  of  dissociation  corresponding  to  a 
particular  temperature.  This  may  be  explained  on  the  basis  of  the 
kinetic  molecular  theory.  Let  us  imagine  the  gas  nitrogen  per- 
oxide to  be  at  a  temperature  below  that  at  which  dissociation 
begins,  when  all  the  molecules  will  have  the  composition  N2O4. 
The  molecules  of  the  gas  are  in  a  state  of  rapid  movement,  and  the 
rapidity  of  their  movement  is  increased  by  rise  of  temperature. 
But  the  molecules  in  a  given  volume  of  the  gas  do  not  all  move 
at  the  same  velocity,  and  therefore  they  have  not  all  the  same 
temperature.  On  account  of  the  infinite  complications  in  their 
movements,  caused  by  their  impacts  against  one  another,  some  will 
be  moving  at  a  speed  considerably  greater  than  that  of  the  average, 
and  will  have  a  temperature  proportionally  higher,  while  others 
again  will  have  a  velocity  and  a  temperature  below  the  average. 


92  Introductory  Outlines 

The  observed  temperature  of  the  gas,  therefore,  is  not  that  of  the 
molecules  having  the  highest  or  the  lowest  velocity  and  tempera- 
ture, but  is  the  average  or  mean  temperature  between,  possibly,  a 
very  wide  range. 

On  the  application  of  heat  to  the  gas,  the  observed  or  mean 
temperature  rises,  but  the  velocity  of  some  of  the  molecules,  and 
consequently  their  temperature,  may  have  been  thereby  raised  to 
the  point  at  which  dissociation  takes  place,  and  they  consequently 
separate  into  the  simpler  molecules.  Let  us  suppose  that  the 
observed  temperature  of  the  nitrogen  peroxide  is  26.7°,  and  that  it 
is  maintained  at  this  point.  Although  this  temperature  may  be 
below  the  dissociation  temperature  of  the  molecules,  it  must  be 
remembered  that  it  only  represents  the  mean  temperature,  and  that 
while  some  of  the  molecules  have  a  lower,  some  also  have  a  higher 
temperature.  As  already  mentioned,  at  the  temperature  of  26.7°, 
20  per  cent,  of  the  molecules  are  dissociated  ;  that  is  to  say,  at 
any  given  instant  one-fifth  of  the  total  number  of  molecules  reach 
a  velocity  which  causes  them  to  break  down  into  the  simpler  NO2 
molecules,  which  themselves  then  take  up  independent  movements. 
If,  in  the  process  of  their  movements,  two  of  these  disunited  mole- 
cules come  into  contact  with  each  other  at  a  moment  when  their 
velocities  are  lower  than  that  at  which  they  dissociated,  they  at 
once  reunite,  so  that  at  the  same  instant  some  are  uniting  and 
others  are  dissociating,  and,  the  two  processes  going  on  equally, 
the  percentage  of  disunited  molecules  at  any  moment  is  the  same, 
although  the  actual  molecules  which  are  dissociated  at  one  point 
of  time  may  not  be  the  identical  ones  that  are  in  this  state  at 
another  time.  Let  us  now  suppose  the  gas  to  be  heated  until  the 
registered  (i.e.  the  mean)  temperature  reaches  60.2°,  and  that  it  be 
maintained  at  this  point.  At  this  higher  temperature  a  much 
larger  proportion  of  the  molecules  will  acquire  a  velocity  at  which 
tliey  are  unable  to  hold  together,  namely,  52.04  per  cent.;  but  the 
remainder,  amounting  to  nearly  one-half,  are  still  at  a  temperature 
below  that  at  which  dissociation  takes  place.  Under  these  altered 
conditions  a  greater  number  of  disunions  and  reunions  takes  place 
during  a  given  interval  of  time,  but  the  numbers  are  equal,  and 
therefore  the  equilibrium  exists.  If  once  more  the  gas  be  further 
heated,  until  the  indicated  temperature  is  140°,  then  it  is  found 
that  the  whole  of  the  N2O4  molecules  have  dissociated  into  NO2 
molecules  ;  that  is  to  say,  when  the  mean  temperature  has  reached 
140°,  then  even  those  molecules  that  are  moving  with  the  slowest 


Balanced  Actions  93 

speed  have  reached  the  temperature  of  dissociation.  It  will  be 
evident  that  the  rate  at  which  the  fraction  of  dissociation  in- 
creases, as  the  temperature  of  a  gas  is  gradually  raised,  will  be 
greatest  when  the  mean  temperature  approaches  the  real  dissocia- 
tion temperature  of  the  gas,  for  the  temperature  of  the  greater 
number  of  the  molecules  will  be  coincident  with,  or  very  closely 
approximating  to,  that  point. 

The  vapour  density  of  nitrogen  peroxide,  if  it  could  be  ascertained 
when  all  the  gaseous  molecules  had  the  composition  N2O4,  would 
be  46  ;  while  that  of  the  gas,  when  entirely  dissociated  into  NO2 
molecules,  is  23.  At  temperatures  between  these  extremes,  the  gas, 
consisting  of  mixtures  of  both  molecules,  will  have  a  density  lying 
between  these  figures,  thus  at  27.6°  and  60.2°  the  density  is  38.3  and 
30.1  (see  Nitrogen  Peroxide,  and  also  Phosphorus  Pentachloride). 

The  effect  of  increased  pressure  upon  a  gas  being  to  diminish 
the  mean  free  path  of  the  molecules,  and  thereby  increase  the 
number  of  molecules  in  a  given  space,  the  number  of  impacts 
between  the  molecules  in  a  given  time  will  be  increased.  If, 
therefore,  while  the  nitrogen  peroxide  is  maintained  at  a  constant 
temperature,  say  62.2°,  the  pressure  be  increased,  the  dissociated 
molecules,  having  shorter  distances  to  travel,  and  making  more 
frequent  impacts  in  a  given  time,  will  unite  more  quickly  than 
others  are  being  disunited,  and  a  fresh  condition  of  equilibrium 
will  be  established  for  any  particular  pressure. 

The  case  of  phosphonium  chloride  already  mentioned  may 
be  referred  to  as  an  illustration.  This  compound  is  completely 
dissociated  into  molecules  of  hydrogen  phosphide,  PH3,  and 
hydrochloric  acid,  below  a  temperature  of  o°.  If,  while  at  this 
temperature,  it  be  subjected  to  pressure,  the  dissociated  molecules 
are  caused  to  unite,  and  at  a  pressure  of  thirteen  atmospheres  the 
union  is  complete,  the  whole  of  the  disunited  molecules  having 
combined  to  form  molecules  of  phosphonium  chloride,  PH4C1. 

If  in  the  process  of  dissociation  one  of  the  products  be  with- 
drawn from  the  sphere  of  action,  then  the  process  may  be  carried 
on  to  completion.  For  example,  in  the  case  of  calcium  carbonate 
already  quoted,  if  this  substance  is  heated  in  such  a  manner  that 
as  fast  as  it  dissociates,  the  gaseous  product,  namely  the  carbon 
dioxide,  is  allowed  to  escape  and  so  pass  away  from  the  sphere  of 
action,  the  change  expressed  by  the  equation 

CaCO3  =  CaO  +  CO2 
will  proceed  until  the  whole  of  the  carbonate  has  been  converted 


94  Introductory  Outlines 

into  oxide.  But  if,  on  the  other  hand,  the  action  is  made  to  take 
place  in  a  closed  vessel,  so  that  the  carbon  dioxide  remains  in 
contact  with  the  lime,  then  the  reverse  action  comes  into  operation, 
namely— 

CaO  +  C02  =  CaC03, 

and  a  condition  is  arrived  at  in  which  the  one  action  proceeds  at 
the  same  rate  as  the  other.  The  pressure  exerted  by  the  carbon 
dioxide  under  these  circumstances  is  spoken  of  as  the  dissocia- 
tion pressure  of  the  calcium  carbonate  for  that  particular  tem- 
perature. 

If,  now,  when  this  condition  of  equilibrium  is  established  the 
temperature  be  raised,  the  balance  will  be  disturbed,  and  the 
materials  will  readjust  themselves  to  a  fresh  condition  of  equilibrium 
at  the  higher  temperature  in  which  the  dissociation  pressure  will 
also  be  greater.  For  any  given  temperature,  therefore,  the  dis- 
sociation pressure  is  the  only  possible  pressure  at  which  a  state  of 
equilibrium  can  be  established  between  carbon  dioxide,  calcium 
carbonate,  and  calcium  oxide  ;  for  if  while  the  temperature  is  con- 
stant the  pressure  upon  the  gas  were  to  be  increased  by  external 
means  and  maintained  at  a  higher  point,  union  between  the  carbon 
dioxide  and  lime  would  proceed  until  the  whole  of  the  lime  was 
converted  into  the  carbonate.  On  the  other  hand,  if  the  pressure 
were  to  be  reduced  and  maintained  at  a  lower  point,  then  dis- 
sociation would  go  on  until  the  action  was  complete  and  once 
more  one  of  the  three  interacting  substances  would  cease  to 
exist. 

Increasing  and  diminishing  the  pressure  upon  a  gas  is  obviously 
synonymous  with  increasing  and  diminishing  the  number  of  mole- 
cules in  a  given  volume.  This  in  modern  phraseology  is  called 
the  molecular  concentration  of  the  gas,  which  embodies  the  same 
idea  as  the  expression  active  mass.  From  the  above  illustration, 
therefore,  it  will  be  clear  that  there  is  some  connection  between 
the  molecular  concentration  (or  active  mass)  of  the  carbon  dioxide 
and  the  rate  of  the  chemical  actions  in  question.  This  connection 
is  thus  formulated  (Guldberg  and  Waage)  :  the  rate  of  chemical 
action  is  proportional  to  the  active  mass  (molecular  concentration) 
of  each  of  the  reacting  substances.  Advantage  is  sometimes 
taken  of  these  facts  in  determining  the  vapour-density  of  a  sub- 
stance which  when  heated  dissociates  into  two  gaseous  con- 
stituents. For  example,  phosphorus  pentachloride  when  heated 


Balanced  Actions  95 

dissociates  into  phosphorus  trichloride  and  chlorine  (see  page  466), 
according  to  the  equation — 

PC15  £:  PC13  +  C12. 

But  if  the  active  mass  of  either  the  chlorine  or  the  trichloride  be 
increased  by  adding  more  molecules  of  either  one  of  these  sub- 
stances from  some  other  source,  the  extent  to  which  dissociation 
takes  place  will  be  proportionally  diminished.  Hence,  by  heating 
the  pentachloride  in  an  atmosphere  of  chlorine  and  thereby  greatly 
increasing  the  molecular  concentration  of  this  gas,  dissociation  may 
be  so  far  prevented  that  the  density  of  the  vapour  is  found  to  have 
practically  the  normal  value  for  the  compound  PC16. 


CHAPTER  XI 
ELECTROLYSIS  AND  ELECTROLYTIC  DISSOCIATION 

IF  a  strip  of  pure  zinc  and  a  strip  of  platinum  be  together  dipped 
into  a  vessel  containing  dilute  sulphuric  acid,  neither  metal  is 
affected  by  the  acid,  so  long  as  the  metals  do  not  touch  each  other. 
If  the  ends  of  the  strips  outside  the  liquid  be  joined  by  means  of  a 
metal  wire,  the  zinc  gradually  dissolves  in  the  acid,  and  bubbles 
of  hydrogen  are  disengaged  from  the  liquid  in  contact  with  the 
surface  of  the  platinum  plate  (which  itself  is  otherwise  unaffected 
by  the  acid),  and  at  the  same  time  an  electric  current  passes 
through  the  wire.  So  long  as  the  chemical  action  of  the  sulphuric 
acid  upon  the  zinc  proceeds,  so  long  will  the  electric  current  con- 
tinue to  pass  ;  in  other  words,  chemical  energy  will  be  transformed 
into  electrical  energy.  If  the  wire  be  severed,  the  electric  current 
can  no  longer  pass,  and  the  chemical  action  at  once  stops. 

Such  an  arrangement  constitutes  a  galvanic  or  voltaic  element 
or  cell,  and  a  series  of  such  cells  forms  a  galvanic  battery.  The 
zinc  plate,  or  the  end  of  a  wire  that  may  be  connected  to  it,  is 
termed  the  negative  pole  of  the  battery,  while  the  end  of  a  wire 
attached  to  the  platinum  plate  is  the  positive  pole.  Other  arrange- 
ments can  be  employed  for  generating  a  galvanic  current,  but  in 
all  cases  the  electrical  energy  is  derived  ultimately  from  chemical 
action. 

If  the  two  poles  of  a  battery  are  connected  together  by  placing 
them  both  in  contact  with  various  different  substances,  it  is  seen 
that  in  some  cases  the  electric  current  passes,  and  in  others  not. 
For  instance,  if  the  poles  are  joined  by  placing  them  both  in  contact 
with  a  bar  of  sulphur,  no  current  passes,  whereas  when  connected 
by  a  rod  of  graphite  the  current  freely  passes.  Substances  which 
behave  in  this  respect  like  the  sulphur  are  said  to  be  non-con- 
ductors of  electricity,  while  those  that  allow  the  current  to  pass 
are  distinguished  as  conductors.  Substances  capable  of  conducting 

electricity  are  of  two  kinds,  namely,  those  which  are  merely  heated, 

96 


Electrolysis  and  Electrolytic  Dissociation  97 

and  those  which  undergo  a  chemical  change  in  consequence.  All 
the  metals,  and  a  few  of  the  non-metals,  belong  to  the  first  of  these 
classes  ;  while  the  second  includes  a  large  number  of  compound 
substances,  which  are  either  in  the  liquid  state  or  in  solution  in 
some  solvent.  Thus,  if  the  poles  of  a  battery  are  immersed  in  pure 
water,  practically  no  current  passes,  because  this  liquid  is  a  non-con- 
ductor ;  but  if  a  quantity  of  hydrochloric  acid  (HC1)  be  dissolved  in 
the  water,  the  solution  at  once  becomes  a  conductor,  and  it  is  seen 
that  gas  is  disengaged  from  the  liquid  upon  the  surface  of  each 
wire.  If  the  solution  of  hydrochloric  acid  is  moderately  strong, 
it  will  be  found,  upon  examination,  that  the  gas  evolved  at  the 
negative  pole  is  hydrogen,  while  that  from  the  positive  pole  is 
chlorine  :  the  hydrochloric  acid,  therefore,  is  separated  into  its 
elements  by  the  passage  of  an  electric  current  through  its 
aqueous  solution.  Such  a  process  is  termed  electrolysis;  and 
the  conducting  liquid  is  known  as  an  electrolyte. 

The  poles  or  terminals  that  are  introduced  into  the  electrolyte 
are  called  electrodes,  the  negative  electrode  being  termed  the 
cathode,  and  the  positive  electrode  the  anode. 

Liquids  which  do  not  conduct  electricity,  or  conduct  only  with 
extreme  difficulty,  such  as  water,  benzene,  aqueous  solutions  of 
alcohol  or  of  sugar,  are  called  non-electrolytes',  while  those  which 
are  good  conductors,  such  as  aqueous  solutions  of  hydrochloric 
acid  or  of  sodium  chloride,  are  called  electrolytes.  Other  liquids 
range  themselves  between  these  two  extremes  with  respect  to 
their  conductivity,  but  those  which  may  be  said  to  fall  about 
midway  are  sometimes  spoken  of  as  half-electrolytes.  These 
terms,  strictly  speaking,  apply  to  the  actual  liquids  or  solutions  ; 
thus  in  the  above  examples  it  is  the  aqueous  solution  of  sugar 
which  is  the  non-electrolyte,  and  the  aqueous  solution  of  sodium 
chloride  which  is  the  electrolyte.  For  brevity,  however,  it  is  usual 
to  apply  the  terms  to  the  substance  in  solution,  and  to  understand 
that  an  aqueous  solution  is  meant  unless  another  solvent  is  specially 
mentioned.  Thus,  when  we  say  that  sugar  is  a  non-electrolyte, 
and  sodium  chloride  an  electrolyte,  it  is  the  aqueous  solutions  of 
these  substances  that  are  referred  to. 

In  the  class  of  electrolytes  are  included  the  strong  acids,  such 
as  nitric,  hydrochloric,  and  sulphuric  acids  ;  the  strong  bases,  such 
as  the  hydroxides  of  the  alkali  metals,  and  almost  all  the  class  of 
substances  known  as  salts,  irrespective  of  whether  the  acids  and 
bases  they  are  composed  of  are  electrolytes  or  half-electrolytes. 

G 


98  Introductory  Outlines 

The  half-electrolytes  are  the  weak  acids,  such  as  acetic,  tartaric, 
and  oxalic  acids,  and  the  weak  bases,  as  ammonium  hydroxide  and 
the  hydroxides  of  divalent  metals  other  than  the  alkaline  earth 
metals.  Non-electrolytes  are  substances  of  a  neutral  character 
such  as  sugar,  this  class  including  the  large  majority  of  organic 
compounds  which  do  not  happen  to  fall  under  the  category  of 
acids,  bases,  and  salts. 

In  a  great  number  of  instances  the  electrolytic  separation  is 
accompanied  by  certain  secondary  reactions,  caused  by  the 
action  of  the  primary  products  of  the  electrolysis  upon  either  the 
electrolyte  or  the  solvent  ;  for  example,  when  a  solution  of  sodium 
chloride  (NaCl)  is  electrolysed,  the  primary  products  are  sodium  and 
chlorine,  the  latter  appearing  at  the  anode  and  the  sodium  making 
its  appearance  at  the  cathode.  The  sodium,  however,  in  contact 
with  the  water  in  the  neighbourhood  of  the  cathode  at  once  reacts 
with  the  liquid,  with  the  liberation  of  its  equivalent  of  hydrogen, 
according  to  the  equation  — 


Similarly,  in  the  case  of  hydrochloric  acid,  if  the  solution  is 
sufficiently  dilute  the  final  products  obtained  by  subjecting  it  to 
electrolysis  are  not  hydrogen  and  chlorine,  but  hydrogen  and 
oxygen.  The  primary  products  are  the  same  as  before,  but  under 
the  altered  condition  the  chlorine  which  is  discharged  at  the 
anode  acts  upon  the  water,  combining  with  the  hydrogen,  and 
liberating  an  equivalent  quantity  of  oxygen  :  the  two  actions 
being  expressed  by  the  equations  — 

4HC1  =  2C1,  +  2H 


Again,  when  a  dilute  solution  of  sulphuric  acid  in  water  is 
electrolysed,  the  acid  separates  into  the  two  primary  products 
H2  and  SO4.  The  hydrogen  as  before  appears  at  the  cathode, 
while  the  group  or  radical  SO4  passes  to  the  anode,  where  it  under- 
goes decomposition  in  contact  with  the  water,  reforming  sulphuric 
acid,  while  oxygen  escapes.  Thus  — 

2H2S04  =  2 


It  will  be  observed  that  the  final  products  are  oxygen  and 
hydrogen  in  the  proportion  of  two  -volumes  of  hydrogen  to  one 


Electrolysis  and  Electrolytic  Dissociation          99 

volume  of  oxygen ;  that  is,  the  proportion  in  which  they  exist 
in  water.  This  process  is,  in  fact,  the  same  as  that  frequently 
spoken  of  as  the  "  electrolysis  of  water." 

If  instead  of  a  solution  of  sulphuric  acid,  a  solution  of  sodium 
sulphate,  Na2SO4,  is  treated  in  the  same  way,  this  compound 
separates  into  the  two  primary  products  2Na  and  SO4 ;  the 
sodium  appearing  at  the  cathode  and  the  SO4  at  the  anode. 
The  sodium  in  contact  with  the  water  reacts  as  explained  above, 
liberating  an  equivalent  quantity  of  hydrogen  ;  while  the  SO4 
group,  as  before,  gives  rise  to  the  reformation  of  sulphuric 
acid  and  the  liberation  of  oxygen.  The  final  products,  there- 
fore, are  again  hydrogen  and  oxygen  in  the  same  proportions 
as  before. 

In  the  same  way,  when  an  aqueous  solution  of  copper 
sulphate  (CuSO4)  is  submitted  to  electrolysis,  the  primary 
products  are  copper,  Cu,  and  the  group  SO4.  The  copper  is 
liberated  at  the  cathode,  and  since  it  exerts  no  action  upon 
the  water,  it  is  deposited  as  a  metallic  film  upon  the  electrode.* 
The  group  SO4  again  passes  to  the  anode,  where  it  undergoes 
decomposition  in  the  presence  of  the  water,  as  in  the  former 
cases,  t 

Faraday's  Law. — When  the  same  quantity  of  electricity  is 
passed  through  different  electrolytes,  the  ratio  between  the 
quantities  of  the  liberated  products  of  the  electrolysis  is  the 
same  as  that  between  their  chemical  equivalents.. 

Thus,  if  the  two  electrolytes,  hydrochloric  acid  and  dilute  sul- 
phuric acid,  be  introduced  into  the  same  electric  circuit,  hydrogen 
and  chlorine  are  evolved  in  the  one  case  and  hydrogen  and  oxygen 
in  the  other.  If  the  gases  be  all  collected  in  separate  measuring 
vessels,  it  will  be  seen  (i)  that  the  hydrogen  and  chlorine  evolved 

*  This  is  the  essence  of  the  process  of  electro-plating.  The  metal  to  be  de- 
posited, whether  it  be  gold,  silver,  or  nickel,  &c. ,  in  the  form  of  a  suitable  salt 
{usually  a  double  cyanide)  in  aqueous  solution,  forms  the  electrolyte.  The  object 
to  be  plated  is  made  the  cathode,  that  is,  it  is  suspended  in  the  liquid  and  is 
connected  to  the  negative  electrode  of  a  suitable  battery. .  The  anode  consists 
of  a  strip  of  the  metal  to  be  deposited.  Thus  in  silver  plating,  a  strip  of  silver 
is  employed,  and  in  this  way  the  acidic  radical  that  is  liberated  at  the  anode 
dissolves  the  metal,  and  thereby  prevents  the  weakening  of  th^  solution, 
which  would  otherwise  result  from  the  gradual  deposition  of  silver  upon  the 
cathode. 

t  For  fuller  explanation  of  these  changes  see  page  207. 


IOO  Introductory  Outlines 

from  the  hydrochloric  acid  are  equal  in  volume  ;  (2)  that  the 
volume  of  hydrogen  collected  from  the  other  electrolyte  is  the  same, 
while  that  of  the  oxygen  is  equal  to  only  one-half  this  amount. 
Knowing  the  relative  weights  of  equal  volumes  of  these  three  gases 
to  be  hydrogen,  oxygen,  chlorine,  as  i,  16,  35.5,  we  see  that  they 
must  have  been  liberated  in  the  proportions  by  weight  of — 

Hydrogen  =  I  Oxygen  =  8  Chlorine  =  35.5. 

Similarly,  if  the  same  quantity  of  electricity  be  passed  through 
aqueous  solutions  of  hydrochloric  acid  (HC1),  silver  nitrate  (AgNO3), 
copper  sulphate  (CuSO4),  and  gold  chloride  (AuCl3),  by  the  time 
that  I  gramme  of  hydrogen  has  been  liberated  from  the  hydro- 
chloric acid,  there  will  be  deposited  upon  the  cathodes  of  ,the  other 
electrolytic  cells  108  grammes  of  silver,  31.7  grammes  of  copper, 
and  65.6  grammes  of  gold.  These  numbers,  which  are  the  electro- 
chemical equivalents,  are  identical  with  the  chemical  equivalents  of 
those  elements,  the  chemical  equivalent  of  an  element  being  its 
atomic  weight  divided  by  its  valency. 

H.  O.  Cl.  Ag.  Cu.  Au. 

Atomic  weights .     .     I         16        35.5         108         63.5         197 

Valency     .     .     .     .     I  2  I  I  2  3 

Regarding  the  quantity  of  electricity  required  to  liberate  r 
gramme  of  hydrogen  as  the  unit,  we  may  say  that  16  grammes  ot 
oxygen  require  2  units  of  electricity  for  its  liberation,  108  grammes 
of  silver  I  unit,  63.5  grammes  of  copper  2  units,  and  197  grammes 
of  gold  3  units  ;  or,  in  other  words,  the  number  of  units  of 
electricity  required  to  liberate  a  gramme-atom  is  identical  with 
the  number  representing  the  valency  of  that  atom  in  the  particular 
electrolyte  employed. 

Some  metals,  such  as  copper,  mercury,  tin,  &c.,  are  capable  of 
functioning  with  different  degrees  of  valency.  Thus  copper  is 
divalent  in  copper  sulphate  and  in  cupric  chloride,  but  mono- 
valent  in  cuprous  chloride.  If,  therefore,  i  unit  of  electricity  be 
passed  through  aqueous  solutions  of  each  of  these  copper  chlorides, 

in  the  case  of  cupric  chloride  —  --  =  31.7  grammes  of  copper  will 

be  deposited,  while  in  the  cuprous  chloride  - —   =  63.5   grammes 

are  formed. 

The  Ionic  Theory. — The  modern  theory  now  generally  held, 


The  Ionic  Theory  101 

to  explain  the  phenomena  of  electrolysis,  is  known  as  the  theory 
of  electrolytic  dissociation  or  the  ionic  theory.  The  passage  of 
electricity  through  conductors  of  the  two  classes  above  mentioned, 
that  is,  through  conductors  such  as  metals,  and  those  which  are 
electrolytes,  may  be  compared  with  the  two  ways  by  which  heat 
is  transmitted,  namely,  by  conduction  and  convection.  When 
a  bar  of  metal  is  heated  at  one  end,  the  heat  travels  along 
the  bar,  the  metal  remaining  stationary ;  but  when  water  is 
contained  in  a  tube  which  is  heated  at  its  lower  end,  the  heated 
particles  of  water  travel  along  the  tube,  conveying  the  heat 
to  the  other  extremity.  In  a  similar  manner,  when  electricity 
passes  through  a  metallic  conductor,  the  electricity  travels  through, 
or  along,  the  metal,  which  itself  does  not  move  ;  *  but  when 
it  is  passed  through  an-  electrolyte,  it  is  conveyed  or  transported 
through  the  liquid  by  the  moving  particles,  to  which  the  name 
ions  (signifying  wanderers)  was  first  given  by  Faraday.  One  set 
of  ions  charged  with  negative  electricity  travels  towards  the  anode, 
while  another  set  conveying  positive  electricity  moves  towards  the 
cathode.  Inasmuch  as  the  negative  ions  appear  at  the  anode  they 
are  called  anions,  while  the  positively  charged  ions  are  distinguished 
as  cations.  In  the  earlier  stages  of  the  development  of  the  present 
theory  it  was  supposed  that  the  electrolyte  was  only  separated 
into  its  ions  as  the  electric  current  was  passed  into  it,  that  the 
electricity  was  the  prime  cause  of  the  dissociation  of  the  electro- 
lyte, hence  the  expression  electrolytic  decomposition^  still  commonly 
used.  It  was  believed  (Grotthus)  that  the  first  effect  of  the  current 
was  to  cause  the  molecules  in  the  solution  to  take  up  positions 
towards  each  other  and  the  electrodes  which  may  be  crudely 
represented  by  the  top  line  in  the  following  diagram,  where  the 
molecules  of  hydrochloric  acid,  for  example,  are  arranged  with 
their  electro-negative  constituents  all  directed  to  the  anode,  and 
their  electro-positive  elements  towards  the  cathode,  precisely  as  a 
number  of  separate  cells  in  a  battery  would  be  arranged.  Then 
that  a  disruption  of  the  molecules  took  place  in  which  those 
nearest  to  the  electrodes  parted  with  their  positive  and  negative 
ions  to  their  respective  electrodes  (where  they  would  be  disengaged 
as  free  hydrogen  and  chlorine  in  the  case  of  hydrochloric  acid), 
while  an  exchange  of  partners  between  the  other  molecules  all 
along  the  line  took  place,  as  represented  in  the  second  line,  result- 
ing in  the  formation  of  fresh  molecules  of  the  original  compound. 

*  In  the  language  of  the  modern  theory  of  the  atomic  nature  of  electricity, 
it  is  the  electrons  which  travel,  while  the  metal  ions  remain  (probably)  stationary. 


IO2  Introductory  Outlines 

These  would  then  immediately  assume  the  position  of  those  in  the 
upper  row.  This  theory,  while  affording  an  explanation  of  many 
of  the  phenomena  connected  with  electrolysis  (such  as  the  fact 
that  the  ions  are  disengaged  only  at  the  surface  of  the  electrodes, 
and  not  in  the  intervening  space  ;  that  the  appearance  of  the 
liberated  ions  takes  place  simultaneously  at  the  two  electrodes, 
however  far  removed  from  each  other,  &c.),  was  not  capable  of 
satisfying  all  the  facts  of  the  case.  It  was  painted  out  (Clausius) 
that  if  the  electric  current  were  the  actual  cause  of  the  separation 
of  the  molecules  into  their  constituent  ions,  this  ought  to  be  made 
manifest  by  the  fact  that  the  current  would  have  to  expend  energy 
in  doing  the  work  of  effecting  such  decomposition.  But  exact 
experiment  shows  that  this  is  not  the  case.  It  is  found  that  when 
an  electric  current  passes  through  an  electrolyte,  no  electric  energy 
is  absorbed  in  causing  the  dissociation  of  the  molecules  of  the 


H  Cl  HCl  HCl  H  Cl  H  Cl  H  Cl 

©0        ©Q        ®Q        ®9        ®0        ®0 
/ 

e 

^ 


©        0®        0®        0®        0®        0©        © 

H  Cl  H  Cl  H  Cl  H  Cl  H  Cl  H  Cl 

FIG.  12. 

dissolved  substance  ;  but  that  the  current  is  conducted  by  electro- 
lytes with  the  same  freedom  as  it  is  by  metallic  conductors.  In 
other  words,  it  has  been  shown  that  Ohm's  law  is  equally  appli- 
cable to  electrolytes  as  it  is  to  metals,  namely,  that  the  current  is 
proportional  to  the  electro-motive  force  for  all  values  of  that  force. 
The  theory  of  electrolytic  dissociation,  first  proposed  by  Arrhenius, 
and  now  generally  accepted  by  chemists  and  physicists,  is  that  all 
solutions  which  are  capable  of  conducting  electricity  contain  mole- 
cules which  are  already  in  a  state  of  dissociation.  That  is  to  say, 
the  electrolyte  consists  of  molecules  which  are  already  dissociated 
into  their  constituent  ions  to  a  greater  or  less  extent.  The  simple 
act  of  solution  in  water  results  in  the  dissociation  of  a  portion  of 
the  molecules  into  their  positive  and  negative  ions.  For  example, 
a  solution  of  sodium  chloride  is  an  electrolyte  ;  when,  therefore, 
this  substance  is  dissolved  in  water  a  certain  proportion  of  the 
molecules  immediately  undergoes  ionic  dissociation,  so  that  the 
solution  contains  some  molecules  of  sodium  chloride,  some  sodium 
ions,  and  some  chlorine  ions  ;  a  state  of  balance  or  equilibrium 
between  the  ions  and  the  undissociated  molecules  being  main- 


The  Ionic  Theory  103 

tained,  depending  upon  various  conditions.  In  such  solutions 
it  is  the  ions  alone  which  take  any  part  in  the  conduction  of 
the  electric  current,  the  undissociated  molecules  being  entirely 
inoperative.  Obviously,  therefore,  when  a  substance  dissolves 
in  water  without  undergoing  ionic  dissociation,  the  solution 
will  be  a  non-electrolyte  ;  while  if  dissociation  only  takes  place 
to  a  limited  extent  the  solution  will  come  under  the  head  of 
the  half-electrolytes.  Strong  acids,  bases,  and  salts,  which  are 
good  electrolytes,  are  therefore  the  substances  which  undergo 
dissociation  to  the  greatest  extent.  For  any  given  solution  the 
extent  to  which  dissociation  takes  place  increases  as  the  solution 
is  diluted  until  a  point  is  reached  at  which  all  the  molecules  are 
dissociated  into  their  ions. 

At  first  it  might  appear  contrary  to  established  ideas  that  in 
such  a  case  as  sodium  chloride,  for  instance,  the  sodium  and 
chlorine  in  the  free  or  separated  state  should  be  capable  of  exist- 
ence side  by  side  in  the  same  liquid — a  liquid,  moreover,  upon  which 
one  of  these  elements,  namely,  the  sodium,  is  under  ordinary  circum- 
stances capable  of  exerting  a  chemical  action.  Similarly,  that  with 
such  a  compound  as  sodium  sulphate  there  should  not  only  be  the 
same  element,  sodium,  existing  in  contact  with  water,  but  also  a 
group  of  elements,  or  radical,  SO4,  which  is  not  known  in  a  state  of 
separate  existence.  These  ions,  however,  whether  elementary  like 
sodium  or  compound  like  the  group  SO4,  are  all  united  with  and 
carry  with  them  enormous  electrical  charges,  positive  or  negative> 
as  the  case  may  be  ;  and  it  is  only  so  long  as  they  retain  their 
electrical  charges  that  they  can  retain  an  independent  existence 
and  exhibit  their  own  special  properties.  When  the  electrodes 
from  an  electric  battery  are  introduced  into  a  solution  of  sodium 
chloride,  the  sodium  ions  with  their  positive  charges  are  attracted 
to  the  cathode  ;  they  there  discharge  their  loads  of  electricity,  and 
thereupon  become  ordinary  molecules  of  sodium,  possessing  the 
properties  usually  associated  with  that  metal.  Hence,  since 
ordinary  sodium  cannot  exist  in  contact  with  water,  the  metal 
immediately  upon  its  liberation  at  the  cathode  reacts  upon  the  water 
with  which  it  is  in  contact  in  the  manner  usual  to  sodium.  Similarly, 
the  chlorine  ions  with  the  negative  electric  charges  are  endowed  with 
their  own  characteristic  properties,  which  are  retained  so  long  as 
the  atom  is  united  to  the  electricity.  So  soon  as  it  loses  its  charge, 
which  it  does  when  it  conveys  it  to  the  anode,  the  chloride  ion 
then  becomes  a  chlorine  atom,  two  of  which  immediately  unite, 
forming  a  molecule  of  the  element  possessing  the  ordinary  proper- 


104  Introductory  Outlines 

ties  of  chlorine  gas.  If,  therefore,  we  use  the  term  radical  to 
embrace  single  atoms  as  well  as  groups  of  atoms,  we  may  describe 
an  ion  as  a  radical  united  to  an  electric  charge — a  positive  ion 
being  one  which  carries  positive  electricity,  and  a  negative  ion 
being  a  radical  which  is  united  to  a  negative  charge. 

Indeed,  instead  of  regarding  this  subject  as  one  presenting  a 
new  difficulty  to  the  mind,  we  may  even  trace  an  analogy  between 
it  and  another  set  of  ideas  with  which  we  are  already  quite  fami- 
liar. We  know  that  when  two  elements  enter  into  chemical  union 
with  each  other  they  lose  their  own  characteristic  properties,  and 
that  the  resulting  compound  is  endowed  with  new  and  different 
properties  ;  when  an  atom  of  sodium  combines  with  an  atom  of 
chlorine  the  sodium  no  longer  exhibits  the  properties  of  metallic 
sodium.  Similarly,  when  an  atom  of  sodium  is  combined  with  a 
negative  electric  charge,  the  product  of  the  union,  namely,  the  ion^ 
possesses  properties  differing  from  those  of  metallic  sodium.  The 
exact  "how"  and  "why"  are  equally  mysterious  in  both  cases, 
and  in  neither  case  are  we  able  to  explain  the  precise  nature 
of  the  union  for  which  in  both  instances  we  employ  the  word 
"combine."  Since  the  immediate  effect  of  passing  an  electric 
current  through  an  electrolyte  is  to  cause  the  ions  to  travel  to 
their  respective  electrodes,  and  there  becoming  electrically  dis- 
charged to  cease  to  exist  as  ions,  it  will  be  evident  that  the 
condition  of  equilibrium  previously  existing  between  the  ions  and 
the  undissociated  molecules  is  at  once  disturbed.  This  disturbance, 
however,  immediately  adjusts  itself  by  the  dissociation  of  more  of 
the  molecules ;  as  fast  as  ions  are  removed  fresh  molecules  dis- 
sociate into  ions.  Hence,  although  the  electric  current  is  not  the 
prime  cause  in  the  production  of  the  ions,  it  is  in  a  sense  an 
indirect  cause,  since  by  bringing  about  the  removal  of  the  ions 
previously  present  it  induces  conditions  which  allow  more  of  the 
molecules  to  dissociate  into  ions. 

Atomic  Electric  Charges — Valency. — If  we  take  as  our  unit 
the  amount  of  electricity  which  is  carried  by  one  atom  of  hydrogen, 
then  of  all  monovalent  ions  we  may  say  that  they  convey  one  unit 
of  electricity,  for  all  such  ions  are  united  to  equal  amounts  of 
electricity,  whether  they  be  simple  or  complex  radicals.  Divalent 
and  trivalent  ions  respectively  are  united  to  two  and  three  units  of 
electricity.  Valency  may,  in  fact,  be  defined  as  the  number  of 
unit  electric  charges  which  are  united  to  an  atom  (or  radical). 
These  electric  charges  are  called  electrons,  or  atoms  of  electricity, 
in  accordance  with  the  present-day  views  as  to  the  nature  of 


The  Ionic  Theory  105 

electricity.  Electricity  is  now  regarded  as  having  an  atomic 
structure  :  it  is  believed  to  consist  of  indivisible  and  inde- 
structible particles,  positive  electrons  and  negative  electrons, 
comparable  in  a  measure  with  the  atoms  of  monovalent  chemical 
elements.  To  denote  these  electrons,  or  atomic  charges  of  elec- 
tricity, the  symbols  +  and  —  are  employed  ;  they  represent  one 
"atom  of  electricity"  (positive  and  negative  respectively),  just  as 
the  symbol  H  stands  for  one  atom  of  hydrogen.* 

A  positive  electron  combined  with  a  positive  chemical  atom 
or  radical  gives  rise  to  a  positive  ion,  or  cation;  while  negative 
elements  or  radicals  united  to  negative  electrons  constitute  nega- 
tive ions  or  anions. 

Ionic  Notation. — In  chemical  notation  it  is  usual  to  represent 
ions  by  employing  either  the  ordinary  ®  and  ©  signs,  or  more 

commonly  a  dot  (')  and  dash  ('),  in  conjunction  with  the  chemical 

+ 
symbol  for  the  atom  or  radical.     Thus  Na  or  Na*  signifies  a  sodium 

ion,  and  Cl  or  Cl'  represents  the  chloride  ion. 

The  symbol  Na'  therefore  conveys  the  information  that  the 
sodium  ion  is  a  monovalent  cation  ;  while  Cl'  indicates  that  the 
chloride  ion  is  a  monovalent  anion.  SO4",  in  the  same  way,  stands 
for  the  sulphate  ion,  with  its  two  negative  charges,  and  Fe'"  for 
the  trivalent  ferric  ion  with  its  triple  charge  of  positive  electricity. 
Sodium  chloride  in  solution  would  be  represented  by  the  formula 
Na'Cl',  ferric  chloride  by  Fe"'Cl'3,  potassium  sulphate  by  K'2SO4", 
and  so  on. 

In  the  system  of  nomenclature  of  the  ions  now  generally  adopted,t 
the  names  of  the  cations  are  formed  by  the  addition  of  the  termina- 
tion ion  to  the  stem  of  the  chemical  name  of  the  element  or  radical ; 
thus,  hydrion,  H',  sodion,  Na',  ammonion,  NH'4,  calcion,  Ca", 
zincion,  Zn",  &c. 

When  it  becomes  necessary  to  indicate  the  number  of  unit 
charges  (i.e.  the  valency)  of  the  radical,  Greek  numerals  are  pre- 
fixed to  the  name.  For  example,  diferrion,  Fe"  (the  ions  in  ferrous 
salts),  triferrion,  Fe"*  (the  ions  in  ferric  salts) ;  monocuprion,  Cu*, 
and  dicuprion,  Cu",  for  the  cations  in  cuprous  and  cupric  com- 
pounds respectively. 

*  Negative  electrons  are  known  in  the  free  state.  The  "cathode"  rays 
emitted  from  a  Geissler  vacuum  tube  consist  of  these  negative  electrons,  and 
they  also  form  a  part  of  the  "  radiation  "  emitted  by  the  element  radium  (see 
Appendix).  So  far  positive  electrons  have  not  been  isolated. 

f  First  introduced  by  J.  Walker. 


io6  Inorganic  Chemistry 

In  the  case  of  anions  the  names  are  formed  by  the  use  of  one  of 
the  three  terminations — idion^  anion,  and  osion,  depending  upon 
whether  the  salt  radical  ends  in  ide,  ate,  or  ite.  For  instance, 
anions  derived  from  chlonVfey,  btonuV&r,  hydroxzVfe.y,  sulphzVzky,  will 
be  c\\\ondion  Cl',  bronwV#0#  Br',  hydroxzV/z'0/z  OH',  sulph/dftp*  S" 
respectively ;  those  from  chlora/£f,  sulphcz/^,  orthophosph<7/<?.r, 
&c.,  chlora/zzVw  C1O3',  sulph<2;zz'tf/z  SO4",  orthophosphdTzzVw  PO4'", 
&c.  ;  while  those  derived  from  such  salts  as  nitr//w  and  sulphz'/^ 
are  termed  ri\\xosion  NCV,  sulphosion  SO3".  These  names  are 
employed  precisely  as  ordinary  chemical  names  are  used,  that  is 
to  say,  they  apply  to  the  material  taken  collectively,  and  not  to 
the  particles  themselves  of  which  the  material  is  composed.* 

It  is  often  convenient  to  regard  the  amount  of  electricity  which 
is  carried  by  one  gramme  of  hydrogen  as  the  unit,  instead  of  that 
conveyed  by  one  atom.  The  value  of  this  unit  is  96,550  coulombs. 
Hence  these  dots  and  dashes  signify  that  one,  two,  or  three  times 
96,550  coulombs  of  electricity  are  carried  by  the  gramme-molecule 
(see  p.  57)  of  the  ion  according  to  the  number  of  these  signs 
attached  to  it.  Thus  96  grammes  of  SO4"  will  carry  96,550x2 
coulombs  of  negative  electricity  ;  18  grammes  of  NH4*  carries 
96,550x1  coulombs  of  positive  electricity,  and  95  grammes  of 
PO4'"  conveys  96,550x3  coulombs,  or  3  units  of  electricity.  In 
other  words,  each  dot  and  dash  attached  to  the  formula  signifies 
one  charge  of  96,550  coulombs  united  to  the  gramme-molecule  of 
the  ion. 

*  Just  as  the  names  sodium,  hydrogen,  chlorine,  &c.,  are  used  to  denote 
matter  which  is  made  up  of  atoms  or  molecules  of  sodium  hydrogen  or  chlorine 
respectively,  so  the  terms  sodion,  hydrion,  chloridion,  are  the  names  applied 
to  the  matter  which  is  composed  of  sodium  ions,  hydrogen  ions,  and  chloride 
ions  respectively.  We  speak  of  a  sodium  atom,  and  of  hydrogen  molecules,  so 
also  of  a  sodium  ion  and  hydrogen  ions.  But  to  use  such  expressions  as  a 
sodion,  or  hydrions,  is  as  meaningless  as  to  speak  of  a  sodium  or  hydrogens. 

The  translation  of  the  chemical  equation 

2HC1  =  H2  +  C12 

is  that  hydrochloric  acid  is  decomposed  into  hydrogen  and  chlorine — or  that 
two  molecules  of  hydrogen  chloride  yield  one  molecule  of  hydrogen  and  one 
molecule  of  chlorine — similarly  the  ionic  equation 

HC1  =  H-  +  C1' 

signifies  that  on  solution  in  water  hydrochloric  acid  is  ionised  into  hydrion 
and  chloridion — or  that  a  molecule  of  hydrogen  chloride  yields  on  ionisation 
a  hydrogen  ion  and  a  chloride  ion. 


The  Ionic  Theory  107 

It  will  be  evident  that  ionisation  or  electrolytic  dissociation 
is  a  phenomenon  of  a  different  order  from  that  which  takes 
place  when  a  compound  dissociates  under  the  influence  of  heat, 
as  discussed  in  the  previous  chapter.  Under  these  circumstances 
it  was  explained  that  the  salt  ammonium  chloride,  for  example, 
dissociates  when  heated  into  the  two  compounds  NH3  and  HC1 ; 
whereas  when  it  is  dissolved  in  water  it  undergoes  electrolytic 
dissociation  into  the  two  ions  NH4*  and  Cl'  ;  in  the  first  case  the 
products  are  electrically  neutral  chemical  compounds,  while  in  the 
latter  they  are  electrically  charged  ions,  or  compounds  of  radicals 
with  electrons. 

From  the  point  of  view  of  the  ionic  theory,  acid,  bases,  and 
salts  all  behave  in  a  perfectly  similar  manner  ;  to  the  "  ionist,"  as 
such,  there  is  no  difference  between  these  three  kinds  of  sub- 
stances ;  it  is  therefore  sometimes  convenient  to  class  them  all 
together  as  salts.  Those  which  from  a  chemical  point  of  view  are 
acids,  from  the  ionic  standpoint  are  salts  of  hydrogen,  that  is,  salts 
in  which  all  the  positive  ions  are  hydrogen  ;  while  those  which 
are  usually  termed  bases  are  spoken  of  as  salts  ofhydroxyl,  or  salts 
in  which  the  only  negative  ions  are  hydroxide  ions.* 

Molecular  Conductivity, — What  is  understood  as  the  molecular 
conductivity  of  a  solution  is  its  specific  conductivity  expressed  in 
the  usual  electrical  units,  divided  into  the  number  of  gramme- 
molecules  of  the  dissolved  substance  contained  in  the  solution  ;  or 
what  is  the  same,  multiplied  by  the  number  of  litres  of  the  solution 
which  contains  one  gramme-molecule  of  the  substance. 

Now  since  it  is  the  ions  present  in  an  electrolyte  which  alone 
take  any  part  in  the  conveyance  of  electricity,  the  undissociated 
molecules  present  being  inoperative,  it  will  be  obvious  that  the 
molecular  conductivity  of  an  electrolyte  will  depend  partly  upon 
the  number  of  ions  present — in  other  words,  upon  the  extent  to 
which  the  electrolyte  is  dissociated — and  partly  upon  the  rate  at 
which  the  ions  travel  or  migrate  in  the  liquid. 

It  has  been  found  (Hittorf)  that  different  ions  under  the  same 
conditions  travel  at  different  rates.  From  determinations  of  the 
changes  in  concentration  which  take  place  in  the  electrolyte 

*  The  student  will  not  fall  into  the  error  of  supposing  that  it  would  be 
either  desirable  or  possible  to  abolish  the  classification  of  acids,  bases,  and 
salts.  From  a  purely  chemical  standpoint  acids  and  bases  are  two  perfectly 
distinct  classes  of  compounds,  and  these  two  terms  will  always  be  employed 
to  denote  them. 


io8  Introductory  Outlines 

immediately  round  the  electrodes,  it  has  been  shown  that  in  a 
solution  of  given  concentration  and  under  the  same  electrical 
conditions,  all  the  ions  of  one  kind  travel  with  a  constant  velocity, 
but  that  the  rate  differs  for  different  kinds  of  ions.  For  example, 
it  is  found  that  the  ion  H*  migrates  with  a  velocity  about  twice  as 
great  as  that  at  which  the  negative  ion  HO'  travels,  and  about  five 
times  the  rate  at  which  the  cation  K'  migrates. 

When,  therefore,  a  solution  is  diluted,  and  its  molecular  con- 
ductivity thereby  increased,  this  increased  conductivity  will  be  due 
partly  to  the  greater  rate  of  migration  of  the  ions  which  follows 
upon  dilution,  and  partly  to  the  increased  number  of  ions  present ; 
for,  as  already  stated,  as  the  solution  is  diluted  more  and  more,  so 
ionisation  takes  place  to  a  greater  extent. 

It  is  found  by  experiment  that  as  the  solution  is  diluted,  the 
molecular  conductivity  at  first  rises  somewhat  rapidly,  that  is  to 
say,  a  moderate  increase  of  dilution  causes  a  considerable  rise  in 
conductivity ;  but  after  a  certain  dilution  is  reached,  the  rate  of 
increase  of  molecular  conductivity  is  greatly  diminished;  and 
after  continuing  slowly  to  increase  on  further  dilution,  a  point  is 
at  length  reached  beyond  which  no  increase  of  conductivity  follows 
upon  additional  dilution.  The  conductivity  at  this  latter  point  is 
called  the  molecular  conductivity  at  infinite  dilution,  and  at  this 
point  the  whole  of  the  electrolyte  has  become  dissociated  into  its 
ions.  The  point  of  dilution  at  which  the  rate  of  increase  of  mole- 
cular conductivity  makes  the  marked  change  may  be  regarded  as 
the  point  at  which  dilution  ceases  to  influence  the  rate  of  migra- 
tion of  the  ions. 

Since  the  molecular  conductivity  is  in  this  way  dependent  upon 
two  factors,  namely,  the  speed  of  migration  of  the  ions  and  the 
degree  of  ionic  dissociation,  it  will  be  obvious  that  it  cannot  by 
itself  afford  a  true  measure  of  dissociation.  The  dissociation 
coefficient,  or  the  fraction  of  the  molecules  of  an  electrolyte  which 
are  dissociated  into  their  ions  at  a  given  concentration,  is  the  ratio 
between  the  molecular  conductivity  at  that  concentration  to  the 
molecular  conductivity  at  infinite  dilution.  Hence,  if  in^  and 
;«c  are  the  molecular  conductivities  at  definite  dilution  and  at 
concentration  c  respectively,  then  the  coefficient  of  dissociation  d 
will  be — 

d^-mclm^. 

Some  general  idea  of  the  degrees  of  dilution  which  are  being 
dealt  with  in  these  considerations  may  be  gained  from  a  single 


TJie  Ionic  Theory  109 

example.  Thus  in  a  solution  of  common  salt,  the  strength  of  the 
solution  at  which  the  rate  of  the  migration  of  the  ions  is  practically 
unaffected  by  further  dilution  is  such  that  one  litre  contains 
about  y^th  of  a  gramme-molecule  of  the  salt,  or  5.85  grammes  ; 
while  a  solution  which  has  been  diluted  until  its  molecules  are 
wholly  dissociated  contains  only  about  -foono^  °^  a  gramme- 
molecule  per  litre,  or  is  a  thousand  times  more  dilute. 

Some  Applications  of  the  Ionic  Theory.  —  The  ionic  theory  is  in 
harmony  with  and  derives  support  from  the  laws  which  regulate 
the  influence  of  substances  in  solution  upon  osmotic  pressure 
(page  158),  upon  the  lowering  of  the  vapour-pressure  (page  135), 
and  upon  the  lowering  of  the  freezing-point  of  the  solvent 
(page  140).  Dilute  solutions  of  electrolytes  are  found  to  ex- 
hibit deviations  from  these  laws  much  in  the  same  way  that 
gases  which  undergo  dissociation  depart  from  the  usual  gas  laws. 
Thus  it  is  observed  that  in  the  case  of  dilute  solutions  of  electro- 
lytes, the  osmotic  pressure,  the  lowering  of  the  vapour-pressure, 
and  the  lowering  of  the  freezing-point  of  the  solvent,  instead  of 
being  proportional  to  the  number  of  molecules  of  the  dissolved 
substance,  are  proportional  to  the  number  of  dissociated  ions. 

Again,  this  theory  affords  an  explanation  of  the  fact  that  the 
heat  of  neutralisation  of  one  equivalent  of  strong  acids  and  bases 
(in  dilute  solution)  is  practically  a  constant,  namely,  about  13,700 
heat  units  or  calories  (see  page  165).  Now,  in  the  neutralisation 
of,  say,  nitric  acid  by  potassium  hydroxide,  according  to  the  ionic 
theory  these  two  reacting  substances  are  in  a  state  of  dissociation 
in  the  dilute  solution  ;  moreover,  the  salt  potassium  nitrate,  result- 
ing from  the  interaction,  will  also  be  dissociated.  The  only  product 
of  the  chemical  action  which  is  not  dissociated  is  the  water,  as  this 
compound  is  practically  a  non-electrolyte  ;*  hence  the  process  of 
neutralisation  of  this  acid  with  this  base  resolves  itself  into  the 
union  of  H'  ions  with  HO'  ions  to  form  molecules  of  H2O,  as  may 
be  seen  by  the  equation 


in  which  the  formulae  for  the  dissociated  molecules  are  written  with 
their  ions  separated  by  a  comma.     It  will  be  obvious,  therefore, 

• 

*  Probably  there  is  no  such  thing  as  an  absolutely  perfect  non-electrolyte. 
In  reality  water  itself  undergoes  ionic  dissociation  to  a  very  slight  extent.  It 
has  been  estimated  that  in  ten  million  litres  of  water  there  will  be  about  one 
gramme-molecule  in  the  ionic  state. 


HO  Introductory  Outlines 

that  the  final  result,  namely,  the  union  of  H'  with  HO',  will  be  the 
same  if  we  substitute  other  strong  acids  or  bases,  thus  — 


Therefore  the  heat  of  neutralisation  of  dilute  solutions  of  these 
acids  and  bases  is  in  reality  the  heat  of  formation  of  H2O  mole- 
cules by  the  union  of  H'  ions  with  HO'  ions. 

Similarly,  the  ordinary  "reactions"  employed  in  chemical  analysis, 
when  considered  from  the  standpoint  of  the  ionic  theory,  become 
invested  with  a  new  meaning,  and  are  often  rendered  more  intel- 
ligible :  one  or  two  examples  may  be  given.  When  the  metal  tin 
is  precipitated  from  a  solution  of  stannous  chloride  by  means  of 
metallic  zinc,  the  following  ionic  equation  expresses  the  change  :  — 

Sn",Cl',Cl'  +  Zn  =  Sn  +  Zn",Cl',Cl'. 

In  other  words,  the  two  unit  charges  of  positive  electricity  have 
been  discharged  by  the  tin  ion,  which  then  ceases  to  be  an  ion, 
but  appears  as  ordinary  metallic  tin,  and  are  transferred  to  the 
metal  zinc,  which  then  ceases  to  be  ordinary  metallic  zinc,  but 
passes  into  the  solution  as  a  zinc  ion. 

Again,  the  tests  for  iron  in  the  ferric  state  are  really  tests  for 
triferrion  Fe"',  and  tests  for  this  metal  in  the  ferrous  condition 
are  tests  for  diferrion  Fe".  But  if  a  compound  containing  this 
metal  should  dissociate  in  such  a  manner  as  to  afford  neither 
Fe"*  nor  Fe"  ions,  it  will  be  evident  that  the  usual  reagents  em- 
ployed to  detect  these  ions  will  yield  no  result.  The  salt  potassium 
ferro-cyanide,  K4Fe(CN)6,  is  a  case  in  point.  On  solution  this 
compound  dissociates  into  the  ions  K'  and  Fe(CN)Giv,  and  the  iron 
in  this  solution,  therefore,  does  not  respond  to  the  usual  tests  for 
either  triferrion  or  diferrion. 

Again,  the  action  of  ammonium  chloride  in  preventing  the  pre- 
cipitation of  magnesium  as  hydroxide  by  ammonia,  is  explained  by 
the  fact  that  ammonium  hydroxide  being  a  comparatively  weak 
base  undergoes  dissociation  to  only  a  slight  extent  into  ammonion 
NH4'  and  hydroxidion  OH'  —  to  an  extent  far  smaller  than  is  the 
case  with  sodium  and  potassium  hydroxides.  Upon  the  addition 
of  an  ammonium  salt  of  a  strong  acid,  such  as  hydrochloric  acid, 
we  are  throwing  into  the  solution  a  large  number  of  ammonium 
ions,  which  has  the  effect  of  causing  the  re-union  of  the  hydroxide 
ions  until  practically  the  whole  of  the  ammonium  hydroxide  present 


The  Ionic  Theory  1  1  1 

is  in  the  undissociated  state,  and  as  there  is  now  no  hydroxidion 
present  no  magnesium  hydroxide  can  be  formed. 

When  no  ammonium  chloride  is  added,  partial  precipitation  of 
magnesium  hydroxide  result 


— 

Mg",Cl',Cl'+NH-4,OH'-fNH-4,OH' 
=  Mg(HO)2  +  NH-4,Cl'-f-NH-4,Cl'. 


But  this  process  results  in  the  introduction  into  the  solution  of 
NH*4  ions,  and  equilibrium  is  established  when  these  are  present 
in  sufficient  quantity  to  prevent  further  production  of  hydroxidion 
by  the  dissociation  of  any  more  of  the  ammonium  hydroxide. 

Similarly  the  behaviour  of  many  salts  in  yielding,  when  dissolved 
in  water,  solutions  which  are  either  acid  or  alkaline,  admits  of  an 
ionic  explanation.  Sodium  nitrite  may  serve  as  an  example. 

Unlike  sodium  nitrate,  which  yields  a  neutral  solution,  this  salt 
when  dissolved  in  water  gives  a  solution  which  is  alkaline,  that  is, 
a  solution  containing  hydroxidion.  When  dissolved,  the  salt  is 
first  largely  ionised  into  sodion  and  nitrosion, 

NaNO2  =  Na',NO2'. 

Besides  these  ions,  however,  there  are  also  present  minute  quantities 
of  hydrion  H'  and  hydroxidion  OH'  due  to  the  very  slight  ionisation 
of  the  water  itself,  hence  we  have  the  ions 


Since  nitrous  acid  is  a  weak  acid,  i.e.  one  which  is  only  slightly 
ionised  in  solution,  the  NO2'  and  the  H'  ions  tend  to  unite  to  form 
molecules  of  undissociated  nitrous  acid,  HNO2,  thereby  causing 
more  water  molecules  to  become  ionised,  with  the  consequent 
increase  in  the  number  of  hydroxide  ions  present  This  process 
goes  on  until  equilibrium  is  established,  which  may  be  thus 
represented  — 


Sodium  hydroxide  being  a  strong  base,  the  hydroxidion  and  sodion 
do  not  unite  to  form  molecules,  but  remain  in  the  ionic  state. 

Processes  of  this  order  are  spoken  of  as  hydrolysis  —  the  sodium 
nitrite  in  this  case  is  said  to  be  hydrolysed.  All  salts  of  weak  acids 
with  strong  bases  behave  in  a  similar  manner. 


CHAPTER  XII 
CLASSIFICATION   OF   THE   ELEMENTS 

IT  has  already  been  mentioned  (page  7),  that  the  elements  may 
be  classified  under  the  two  subdivisions,  metals  and  non-metals. 
Further  classifications  have  from  time  to  time  been  in  use,  based 
upon  other  properties,  such,  for  example,  as  the  valency  of  the 
elements. 

Classified  according  to  their  valency,  the  elements  fall  into  six 
subdivisions,  consisting  of  mono-,  di-,  tri-,  tetra-,  penta-,  and  hexa- 
valent  elements.  This  system  of  classification  has  now  largely 
fallen  into  disuse,  owing  partly  to  the  difficulties  arising  out  of  the 
variability  of  valency  so  often  exhibited,  but  more  especially  to  the 
more  recent  development  of  another  system,  known  as  the  natural 
classification  of  the  elements,  or  the  periodic  system^  which  practi- 
cally absorbs  and  includes  the  older  method. 

Certain  remarkable  numerical  relations  have  long  been  observed 
to  exist  among  the  atomic  weights  of  elements  that  closely  re- 
semble one  another  in  their  chemical  habits.  In  such  groups  or 
families  it  is  frequently  seen  that  the  atomic  weight  of  one  mem- 
ber is  approximately  the  arithmetic  mean  of  the  atomic  weights  of 
those  immediately  before  and  after  it,  when  they  are  arranged  in 
order  of  their  atomic  weights.  This  will  be  seen  from  the  following 
examples  : — 

Li.  Na.  K. 

7  23  39 

K.  Rb.  Cs. 

39  85  133 

P.  As.  Sb. 

31  75  I2° 

S.  Se.  Te. 

32  77  125 


The  Periodic  Classification  1 1 3 

If  the  elements  in  these  various  families  are  so  arranged,  as 
to  bring  out  the  differences  between  their  atomic  weights,  the 
striking  fact  will  be  observed  that  the  increase  in  the  atomic 
weights  in  each  group  takes  place  by  practically  the  same  incre- 
ment. In  the  following  table  the  elements  belonging  to  the  same 
group  are  placed  in  vertical  columns,  the  differences  between  the 
various  atomic  weights  being  placed  between  them  : — 


F  =  19 


N=i4 

Difference  .   16.5  j  Diff. 
Cl  =  35-5  P 


O  =  16  Na  =  23 


16    Diff.      . 

:32  K  = 


16 


39 


.      17  Diff.      . 

Difference  .  44.5  Diff.  .  44  |  Diff.  .  47  Diff.  .  46.2 

Br  =  80  As  =  75  Se  =  79  Rb  =  85.2 

Difference  .  47  Diff.  .  45  Diff.  .  46  |  Diff.  .  47.8 

I  =  127  Sb  =  120 


Te  =  125     '     Cs  =  133 


Mg  =  24 
Diff.  .  16 

Ca  =  40 
Diff.  .  47-3 

Sr  =  87.3 
Diff.  .  49.7 

Ba  =  137 


It  will  be  seen  that  in  each  group  the  difference  between  the  first 
and  second  number  is  about  16,  while  between  all  the  others  the 
increase  in  weight  takes  place  by  a  number  which  approximates 
to  16  x  3. 

This  numerical  relation  between  the  atomic  weights  of  elements 
of  the  same  family,  and  between  the  various  groups,  is  obviously 
not  a  chance  one,  and  chemists  were  led  by  it  to  believe  that  the 
properties  of  the  elements  were  in  some  way  related  to  their  atomic 
weights.  Newlands  (1864)  was  the  first  to  point  out,  that  if  the 
elements  are  tabulated  in  the  order  of  increasing  atomic  weights, 
the  properties  belonging  to  each  of  the  first  seven  elements  reap- 
peared in  the  second  seven,  and  he  applied  to  this  relation  the 
name  of  the  law  of  octaves.  A  more  elaborated  and  systematic 
representation  of  Newlands'  law  of  octaves  was  afterwards  deve- 
loped by  Mendelejeff  (1869),  and  which  is  now  generally  known  as 
MendelejefFs  periodic  law.  At  the  present  time,  owing  to  the 
recent  discovery  of  the  argon  family  of  elements,  it  is  not  until 
eight  elements  have  been  traversed  that  the  properties  of  the  first 
reappear  ;  the  term  "  octaves  "  is  therefore  no  longer  strictly 
applicable.* 

*  Unless,  indeed,  we  stretch  the  musical  simile  somewhat  and  look  upon 
these  five  inert  gases  as  "  accidentals." 


114  Introductory  Outlines 

If  the  sixteen  elements  with  lowest  atomic  weights,  after 
hydrogen,  be  arranged  in  order  of  increasing  atomic  weights  in 
two  horizontal  rows  of  eight,  some  of  these  relations  will  be 
recognised — 

He=4     Li  =  7   Be  =9    B  =  11   C  =  i2  N  =  i4  O  =  i6  F=i9. 

Ne  =  2o  Na=23  Mg  =  24  Al  =  27  81  =  28  P  =  3i    8=3201  =  35.5. 

In  traversing  the  upper  row  from  helium  to  fluorine,  we  meet  with 
certain  characteristic  properties  belonging  to  each  member,  and 
also  a  certain  gradation  in  those  properties  that  are  common. 
Coming  to  the  second  row,  many  of  the  characteristic  properties 
of  the  members  of  the  first  row  again  appear,  and  the  same  regular 
modulation  is  met  with  in  passing  along  the  series  :  thus  helium 
exhibits  a  likeness  to  neon,  lithium  resembles  sodium,  carbon 
corresponds  to  silicon,  fluorine  to  chlorine,  and  so  on.  These 
resemblances  are  seen  both  in  the  physical  as  well  as  the  chemical, 
properties  of  the  elements,  thus  lithium  and  sodium  are  both  soft 
white  metals,  and  are  strongly  electro-positive.  Fluorine  and 
chlorine  are  both  pungent  corrosive  gases,  and  are  intensely  electro- 
negative ;  while  helium  and  neon  are  neither  electro-positive  nor 
electro-negative,  have  no  chemical  properties  whatever,  and 
therefore  no  valency.  Taking  their  power  of  combining  with 
chlorine  and  with  hydrogen  as  indicative  of  their  valency,  we  see 
that  the  change  in  this  respect,  as  the  two  series  are  traversed,  is 
the  same  in  each,  thus — 


LiCl      BeCl2      BC13         CC14     CH4      NH3     OH2     FH. 
NaCl     MgCl2     (A1C13)2     SiCl4    SiH4     PH3     SH2     C1H. 

The  gradation  in  properties  exhibited  by  the  elements  in  a  series 
is  also  seen  in  their  power  of  combining  with  oxygen,  which  will 
be  more  clearly  brought  out  if  the  formulas  of  the  compounds  be 
so  written  as  to  indicate  the  relative  proportions  of  oxygen  with 
which  two  atoms  of  each  element  unite,  thus — 

Na20     (Mg20,)     A1203     (Si,O4)     P2O5     (S2OC)     C12O7 
MgO  SiO2  8~O3 

Regarding,  then,  the  eight  elements  of  the  first  row  as  ^period,  we 
find  that  the  various  properties  exhibited  by  the  several  members 
are  met  with  again  in  those  of  the  second  period. 


The  Periodic  Classification  1 1 5 

JN  ot  only  do  the  properties  of  the  elements  themselves  reappear, 
but  also  those  possessed  by  the  various  compounds  they  form  :  thus 
lithium  chloride  (LiCl)  and  sodium  chloride  (NaCl)  strongly  re- 
semble one  another.  The  oxides  of  beryllium  and  magnesium 
(BeO  and  MgO)  have  similar  properties.  The  compounds  of  fluo- 
rine and  chlorine  with  hydrogen  (HF  and  HC1)  closely  resemble 
each  other,  and  so  on. 

This  periodic  reappearance  of  similar  properties,  exhibited  by  the 
elements  and  their  compounds  as  the  atomic  weights  of  the  former 
gradually  increase,  is  thus  stated  by  Mendelejeff  in  his  law  of 
periodicity.  The  properties  of  the  elemetits,  as  well  as  the  proper- 
ties of  their  compounds,  form  a  periodic  function  of  the  atomic 
weights  of  the  elements. 

When  the  tabulation  of  the  elements  according  to  this  system  is 
continued  (after  the  completion  of  the  second  period  with  chlorine), 
it  will  be  seen  that,  beginning  with  argon,  eighteen  elements  have 
to  be  arranged  before  we  meet  with  the  reappearance  of  those  pro- 
perties that  belong  to  the  first  ;  that  is  to  say,  there  are  two 
"  octaves,"  one  containing  eight  members  like  the  former  ones,  and 
one  containing  seven,  and  three  elements  over,  which  in  the  follow- 
ing table  are  placed  within  brackets  : — 


A.* 

K. 

Ca. 

Sc. 

Ti. 

V. 

Cr. 

Mn. 

(Fe. 

Co. 

Ni.) 

40 

39 

40 

44 

48 

51 

52 

55 

(56 

59 

59) 

Cu. 

Zn. 

Ga. 

Ge. 

As. 

Se. 

Br. 

63.5 

65 

70 

72 

75 

79 

80 

This  constitutes  what  is  known  as  a  long  period,  in  contradis- 
tinction to  the  two  first,  which  are  distinguished  as  short  periods. 
In  certain  respects,  however,  the  last  seven  elements  in  this  long 
period  exhibit  resemblances  to  the  seven  in  the  first  portion  (count- 
ing after  the  first  element,  argon)  ;  that  is  to  say,  the  properties 
displayed  by  the  members  of  the  first  period,  which  is  known  as 
the  typical  period,  reappear  twice  over  in  the  long  period.  The 
three  elements  within  the  brackets  are  termed  by  Mendelejeff 
transitional  elements.  Continuing  the  arrangement  from  bromine, 
another  long  period  occurs,  again  containing  three  transitional 
elements  : — 


*  It  will  be  noticed  that  the  element  argon,  A,  is  placed  before  potassium,  K, 
although,  according  to  the  atomic  weights  here  given,  it  would  appear  as 
though  they  should  be  in  the  reverse  order.  This  will  be  discussed  later. 


n6  Introductory  Outlines 

Kr.       Rb.          Sr.          Y.  Zr.         Cb.       Mo.       —          (Ru.  Rh.          Pd.) 

82       85       87.6       89       90.7      93.5     96       ?       (101.7       103       ic6) 

Ag.  Cd.  In.  Sn.  Sb.  Te.  I. 

108         112         114         118         120         125?         127 

It  will  be  seen  that  a  gap  is  left  where  the  eighth  member  of 
the  first  part  of  this  period  should  be,  an  element  which  would 
correspond,  in  this  period,  with  manganese  in  the  period  above. 
This  element  is  at  present  unknown.  The  remaining  elements 
belong  to  three  other  long  periods,  in  which,  however,  the  number 
of  gaps  is  very  considerable,  thus — 

X.  Cs.  Ba.  La.          Ce.         —        -        —        (  —        —        —  ) 

128          133  137  139         140 


Yb. 
172 

—      Ta.    W.   —   (Os.     Ir.     Pt.) 

181   184     (191   193   195) 

Au. 
197 

Hg. 
20O 

Tl.      Pb. 
204     2O7 

Bi.     —      — 
208 

_ 

_ 

Th.    —    Ur. 

—    (  —     _  y 

232        238.5 

Those  elements  that  fall  in  the  first  eight  places  of  the  long 
periods  are  termed  the  even  series,  while  the  last  seven  are  dis- 
tinguished as  the  odd  series;  arranging  them,  therefore,  in  such  a 
manner  as  to  bring  the  odd  and  even  series  into  columns,  we  get 
the  table  on  page  1 1 8. 

In  this  manner  the  elements  are  arranged  in  nine  groups. 
The  first  of  these  groups  contains  the  so-called  "  inert  gases " — 
the  five  new  elements  of  recent  discovery,  which  take  their  place 
rather  outside  this  classification  scheme,  regarding  it  from  a  purely 
chemical  standpoint.  And  as  the  system  of  numbering  the  groups 
of  elements  in  this  periodic  arrangement  has  become  familiarised 
by  long  use,  this  group  containing  the  "inert  gases"  has  been 
numbered  Group  O,  and  the  systematic  numbering  of  the  other 
groups  begins  as  usual.  The  last  group  contains  the  transitional 
elements  that  come  between  the  even  and  odd  series  of  the  long 
periods. 

In  each  of  the  remaining  seven  groups,  the  elements  belonging 


The  Periodic  Classification  117 

to  the  even  series  of  their  respective  long  periods,  are  placed  to  the 
left,  while  those  belonging  to  the  odd  series  are  arranged  on  the 
right-hand  side  of  each  vertical  column.  In  this  way  the  groups  are 
divided  into  the  subdivisions  A  and  B,  in  which  the  resemblance 
between  the  members  is  most  pronounced.  Thus  in  Group  II., 
although  there  are  certain  properties  common  to  all  the  members, 
there  is  a  much  closer  similarity  existing  between  the  elements 
calcium,  strontium,  and  barium  than  between  zinc  and  calcium,  or 
cadmium  and  barium.*  The  elements  in  the  two  short  periods 
have  been  placed  in  that  subdivision  or  family  with  the  members 
of  which  they  exhibit  the  closest  resemblance.  Thus,  in  Group  I. 
lithium  and  sodium  are  more  allied  to  potassium,  rubidium,  and 
caesium,  than  to  copper,  silver,  and  gold;  while  in  Group  VII. 
fluorine  and  chlorine  are  placed  in  the  same  family  with  bromine 
and  iodine,  with  which  they  exhibit  a  close  similarity. 

In  the  eighth  group,  containing  the  transitional  elements,  the 
families  consist  of  the  horizontal  and  not  the  vertical  rows  ;  that  is 
to  say,  the  closest  resemblance  is  between  the  three  transitional 
elements  in  each  series,  elements  whose  atomic  weights,  instead  of 
exhibiting  a  regular  increase,  as  in  the  other  families,  have  almost 
the  same  value,  such  as  Fe  =  56  ;  Co  =  59  ;  Ni  =  59. 

A  glance  at  the  table  shows  that  in  the  last  three  long  periods 
there  is  a  large  number  of  gaps.  It  is  possible  that  these  gaps 
may  represent  elements  which  yet  await  discovery.  This  supposi- 
tion gains  considerable  support  from  the  fact,  that  at  the  time 
Mendelejeff  first  formulated  the  periodic  law,  there  were  three  such 
gaps  in  the  first  long  period,  which  have  since  been  filled  up  by  the 
subsequent  discovery  of  three  new  elements  ;  these  will  be  referred 
to  later. 

The  periodic  recurrence  of  some  of  the  chemical  properties 
is  indicated  in  the  lowest  horizontal  column,  where  the  general 
formulas  of  the  oxygen  compounds  and  the  hydrides  are  given  ;  R 
standing  for  one  atom  of  any  element  in  the  group.  As  explained 
on  page  114,  these  formulae  are  so  written  as  to  show  the  relative 
amount  of  oxygen  to  two  atoms  of  element,  in  order  to  establish 
the  true  relation  between  the  different  groups.  For  example,  the 

*  This,  however,  is  by  no  means  uniformly  the  case  ;  thus  the  element  copper 
(Group  I.)  in  many  of  its  chemical  attributes  is  much  more  closely  allied  to 
mercury  (Group  II.)  than  to  silver;  and  silver,  again,  more  strongly  resembles 
thallium  (Group  III.)  than  either  copper  or  gold,  with  which  it  is  associated 
in  this  system  of  classification. 


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Introductory  Outlines 


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The  Periodic  Classification  119 

oxides  of  the  elements  of  Group  I.  contain  two  atoms  of  the  element 
to  one  of  oxygen,  as  Li2O  ;  but  those  of  the  second  group  only  con- 
tain one  atom  of  the  element,  as  CaO  :  hence  the  general  formula 
is  doubled,  R2O2.  It  will  be  seen,  therefore,  that  the  proportion  of 
oxygen  relative  to  two  atoms  of  the  element  regularly  increases 
from  the  first  group  to  the  eighth.  The  oxides  of  the  members  of 
Group  I.  are  strongly  basic  in  character,  and  in  general  this  basic 
nature  gradually  diminishes  as  we  traverse  the  series,  giving  place 
to  acidic  characteristics,  which  are  strongly  marked  in  the  seventh 
group. 

The  periodic  reappearance  of  the  physical  properties  of  the 
elements  is  seen  in  such  points  as  their  electrical  characters,  their 
malleability,  ductility,  melting-points,  &c.,  all  of  which  are  in 
harmony  with  the  periodic  law  ;  but  in  none  is  it  more  strikingly 
seen  than  in  their  atomic  volumes  in  the  solid  state.  The  atomic 
volumes  of  the  elements  are  the  relative  volumes  occupied  by 
quantities  proportional  to  their  atomic  weights,  or  by  gramme- 
atoms  ;  and  they  are  obtained  by  dividing  the  atomic  weights  of 
the  elements  by  their  specific  gravities.  In  the  case  of  gases,  as 
has  been  already  explained  on  page  40,  the  specific  gravity  is 
the  density  referred  to  hydrogen  as  the  unit  :  the  atomic  volume, 
therefore,  of  such  a  gas  as  oxygen  is — 

^6  =  atomic  weight  _ 
16  =  density 

The  specific  gravities  of  solids  (and  also  liquids)  are  referred  to 
water  as  the  unit,  and  as  I  cubic  centimetre  of  water  weighs 
i  gramme,  the  specific  gravity  of  a  solid  or  liquid  expresses  the 
weight  in  grammes  of  i  cubic  centimetre  of  the  substance.  Dividing 
the  atomic  weight,  expressed  in  grammes,  by  the  weight  in  grammes 
of  i  cubic  centimetre  (/'.<?.  the  specific  gravity),  the  atomic  volume 
will  be  represented  in  cubic  centimetres.  It  must  be  remembered 
that  the  atomic  volumes  do  not  express  the  relative  volumes  that 
are  actually  occupied  by  the  atoms,  they  represent  in  reality  the 
relative  volume  of  the  atoms  plus  the  unknown  volumes  of  the 
spaces  that  separate  them. 

The  following  table  gives  the  specific  gravities  and  the  calculated 
atomic  volumes  of  the  first  and  the  middle  elements  of  the  two 
short  and  two  long  periods,  not  counting  the  group  of  "inert" 
elements  : — 


I2O 


Introductory  Outlines 


Specific 
Gravity. 

Atomic 
Weights. 

Atomic 
Volumes, 

(  Lithium. 

0-59 

7 

II.9 

\  Carbon  .         .                            .         . 

3-0 

12 

4 

,  p    .    ,  j  Sodium  .                                     . 

0.97 

23 

23.6 

(  Silicon    .                                     .         . 

2.4 

28.3 

1  1.8 

AT)    •  A  $  Potassium 
3dl  Iron  (Cobalt-Nickel) 

0.865 
7-79 

39 

56 

45 
7.2 

-tv,  Pm-ir,H  )  Rubidium       .         .                  .         . 
4th          3d  |  (Ruthenium-Rhodium)  Palladium  . 

1-52 
11.4 

106.2 

56.0 
9 

Caesium         

1.88 

133 

70 

From  the  figures  in  the  last  column  it  will  be  seen,  that  beginning 
with  lithium,  11.9,  the  atomic  volume  falls  as  the  middle  element  of 
the  period,  namely,  carbon,  is  reached  ;  after  which  it  again  rises 
and  reaches  a  maximum  with  the  first  member  of  the  second  period, 
namely,  sodium.  In  this  period  the  same  gradual  fall  in  atomic 
volume  is  again  noticed  until  the  middle  element  (silicon)  is 
reached,  when  the  value  of  this  function  of  the  elements  once  more 
rises,  and  a  second  maximum  is  attained  with  the  first  member 
(potassium)  of  the  third  period.  The  two  next  are  long  periods,  and 
the  atomic  volumes  steadily  decrease  until  the  middle  three  (transi- 
tional) elements,  after  which  they  gradually  increase  again  to  a 
maximum  in  rubidium,  the  starting-point  of  the  fourth  period.  In 
the  fourth  period  the  same  thing  once  more  occurs,  the  minimum 
atomic  volumes  being  those  of  the  middle  or  transition  elements, 
after  which  a  maximum  is  again  reached  in  caesium. 

This  periodicity  of  the  atomic  volumes  may  be  graphically 
represented  by  a  curve,  where  the  ordinates  represent  atomic 
volumes  and  the  abscissse  atomic  weights.  This  curve,  which  was 
first  constructed  by  Lothar  Meyer,  is  known  as  Lothar  Meyer's 
curve  (page  121),  and  a  comparison  of  it  with  MendelejefFs  table 
is  most  instructive. 

The  divisions  indicated  by  the  Roman  numerals  correspond  to 
the  different  periods :  Groups  I.  and  II.  being  the  two  short  periods, 
III.  and  IV.  the  two  complete  long  periods,  while  V.,  VI.,  and 
VII.  correspond  to  the  fragmentary  portions  of  the  last  three 
periods. 

The  transitional  elements  of  periods  III.,  IV.,  and  VI.  are  all  to 
be  found  at  the  minima  of  the  large  hollows  ;  separating  the  even 
series  (situated  on  the  descending  portion  of  the  curve)  from  the 
odd  series  which  lie  on  the  ascending  slope.  The  elements  belong- 


The  Periodic  Classification 


121 


122  Introductory  Outlines 

ing  to  the  different  groups  in  Mendelejeff's  table  are  seen  to  occupy 
the  same  relative  positions  upon  the  different  portions  of  this  curve. 

Thus  in  Group  I.  the  elements  Li,  Na,  K,  Rb,  Cs,  are  all  found 
upon  the  maxima  of  the  curve,  and  Cu,  Ag,  and  Au  at  those  points 
at  the  minima  where  the  electro-negative  properties  reappear.  The 
halogen  elements  (chlorine,  bromine,  iodine)  are  seen  in  similar 
positions  upon  the  ascending,  and  the  alkaline  earths  (beryllium, 
magnesium,  calcium,  strontium,  barium)  on  the  descending 
portions. 

When  the  periodic  law  was  first  formulated  by  Mendelejeff 
(1869),  there  were  a  number  of  instances  in  which  the  system  did 
not  harmonise  with  the  then  accepted  atomic  weights  of  the 
elements.  The  discoverer  boldly  asserted  that  the  atomic  weights, 
and  not  the  system,  were  at  fault,  and  in  almost  every  such  case 
the  careful  reinvestigation  of  the  atomic  weights  by  numerous 
chemists  has  proved  the  correctness  of  the  assertion.  One  or  two 
instances  may  be  quoted.  The  element  indium  had  assigned  to  it 
the  atomic  weight  76.  Its  combining  proportion  is  38,  and  being 
regarded  as  a  divalent  element,  its  o'xide  was  believed  to  have 
the  formula  InO.  Having  an  atomic  weight  =  76,  indium  would 
occupy  a  place  between  As  =  75  and  Se  =  79  ;  but  in  the  system 
(see  table  on  page  118)  there  is  no  room  for  an  element  with  such 
an  atomic  weight ;  and,  moreover,  if  indium  be  a  divalent  element 
having  this  atomic  weight,  it  should  come  between  Zn  =  65  and 
Sr  =  87  in  Group  II.,  where  again  there  is  no  room.  Mendelejeff 
made  the  assumption  that  the  oxide  of  indium  had  the  formula 
In2O3,  believing  the  element  to  be  an  analogue  of  aluminium 
(Group  III.).  If  this  be  the  true  composition  of  the  oxide,  the 
atomic  weight  of  the  element  would  be  38  X  3  =  114,  and  indium 
would  then  take  its  place  in  Group  III.,  between  the  elements 
cadmium  =  112  and  Sn  =  118,  in  the  odd  series  of  the  second  long 
period.  Bunsen  afterwards  determined  the  specific  heat  of  indium 
by  means  of  his  ice  calorimeter,  and  found  it  to  be  0.057  : — 

Mean  atomic  heat    Ai_  =  1 12.3  =  atomic  weight  (see  page  48). 
Specific  heat     .     .  0.057 

Hence  114  and  not  76  is  the  accepted  (approximate)  atomic  weight 
of  indium. 

Again,  the  element  beryllium  (formerly  known  as  glucinum)  has 
a  combining  proportion  of  4.6.  Its  chloride  was  believed  to  have 
the  composition  BeCl3,  and  its  oxide  to  be  a  sesquioxide  having 


The  Periodic  Classification  123 

the  formula  Be2O3.  The  atomic  weight  assigned  to  the  element, 
therefore,  was  13.8. 

With  this  atomic  weight  beryllium  would  take  its  place  between 
carbon  =  12  and  nitrogen  =  14;  but  according  to  the  periodic 
classification  there  is  no  room  for  such  an  element,  and,  moreover, 
in  such  a  position  it  would  be  among  elements  with  which  it  has 
no  properties  in  common.  On  the  supposition  that  the  oxide  of 
beryllium  has  the  formula  BeO,  that  is,  that  the  element  is  divalent, 
its  atomic  weight  would  have  to  be  lowered  from  13.8  to  9.1  in 
order  to  maintain  the  same  ratio  between  the  weights  of  metal  and 
oxygen  in  the  compound.  On  this  assumption,  beryllium  would 
fall  into  the  second  place  in  the  first  series,  between  lithium  =  7 
and  boron  =  1 1,  and  in  the  same  group  as  magnesium  and  zinc. 

When  the  specific  heat  of  beryllium  was  determined,  it  gave  the 
value  0.45,  and  this  number  divided  into  the  atomic  heat  constant, 
6.4,  gave  14  as  the  atomic  weight.  In  spite  of  this  evidence  in 
favour  of  the  higher  value  as  the  atomic  weight  of  beryllium, 
Mendelejeff  still  regarded  the  lower  number  as  correct,  and  it 
was  suggested  that  possibly  beryllium,  like  carbon  and  boron 
(elements  also  of  very  low  atomic  weight),  had  an  abnormally 
low  specific  heat  at  ordinary  temperatures.  This  was  found  to  be 
the  case  (see  page  48),  and  at  500°  the  specific  heat  of  beryllium 
was  found  to  be  0.6206.  This  divided  into  6.4  gives  the  value  10 
as  the  atomic  weight,  which  indicates  that  9.1  and  not  13.8  is  in 
reality  the  atomic  weight  of  beryllium. 

Not  only  has  the  periodic  law  been  of  service  in  bringing  about 
the  correction  of  a  number  of  doubtful  atomic  weights,  but  by 
means  of  it  its  originator  was  enabled  to  predict  with  considerable 
certainty  the  existence  of  hitherto  undiscovered  elements,  and 
even  to  predicate  many  of  the  properties  of  these  elements.  As 
already  mentioned,  at  the  time  when  the  periodic  law  was  first 
formulated,  there  were  three  gaps  in  the  system  in  the  first  long 
period,  namely,  No.  4  in  the  even  series  (now  occupied  by  scandium), 
and  Nos.  3  and  4  in  the  odd  series  (now  filled  by  gallium  and 
germanium).  To  the  unknown  elements  which  were  destined  to 
occupy  these  positions,  Mendelejeff  gave  the  names  eka-boron, 
eka- aluminium,  and  eka-silicon  (the  prefix  eka  being  the  Sanscrit 
numeral  one},  and  from  the  known  properties  of  the  neighbouring 
elements  of  the  series  (horizontal  rows  in  the  table,  page  118),  and 
also  of  those  situated  nearest  in  the  same  family  (vertical  columns), 
he  predicted  .some  of  the  prominent  properties  that  would  pro- 


124  Introductory  Oiitlines 

bably  be  possessed  by  these  elements.  Thus  in  the  case  of  eka- 
aluminium,  from  the  known  properties  of  aluminium  and  indium, 
the  neighbouring  elements  in  the  same  family,  and  from  zinc,  the 
contiguous  element  in  the  same  series  (the  4th  place  in  the  series 
being  unoccupied),  MendelejefF  deduced  the  following  properties 
for  the  unknown  element  that  he  called  eka-aluminium  : — 

PREDICTED  PROPERTIES  OF  EKA-ALUMINIUM  (1871). 

(i.)  Should  have  an  atomic  weight  about  69. 

(2. )  Will  have  a  low  melting-point. 

(3.)  Its  specific  gravity  should  be  about  5.9. 

(4.)  Will  not  be  acted  upon  by  the  air. 

(5.)  Will  decompose  water  at  a  red  heat. 

(6. )  Will  give  an  oxide  £72O3,  a  chloride  £72C16,  and  sulphate  £72(SO4)3. 

(7.)  Will  form  a  potassium  alum,  which  will  probably  be  more  soluble  and 
less  easily  crystallisable  than  the  corresponding  aluminium  alum. 

(8. )  The  oxide  should  be  more  easily  reducible  to  the  metal  than  alumina. 
The  metal  will  probably  be  more  volatile  than  aiuminium,  and  therefore  its 
discovery  by  means  of  the  spectroscope  may  be  expected. 

In  the  year  1875  Lecoq  de  Boisbaudran  discovered  a  new 
element  in  a  certain  specimen  of  zinc  blende  (zinc  sulphide),  the 
individuality  of  which  he  first  recognised  by  the  spectroscope, 
the  spectrum  being  characterised  by  a  brilliant  violet  line.  This 
element  he  named  gallium.  The  properties  of  this  metal,  as  they 
were  subsequently  observed,  showed  that  it  was,  in  fact,  the  pre- 
dicted eka-aluminium  of  Mendelejeff,  as  will  at  once  be  seen  by  a 
comparison  of  the  following  facts. 

PROPERTIES  OF  GALLIUM  (discovered  1875). 

(i.)  Atomic  weight— 69.9. 
(2.)  Melting-point,  30.15°. 
(3.)  Specific  gravity,  5.93. 
(4. )  Only  slightly  oxidised  at  a  red  heat. 
(5.)  Decomposes  water  at  high  temperatures. 

(6.)  Gallium  oxide,  Ga2O3.  Gallium  chloride,  Ga2Cl6.  Gallium  sulphate, 
Ga2(S04)3. 

(7. )  Forms  a  well-defined  alum. 

(8.)  Is  easily  obtained  by  the  electrolysis  of  alkaline  solutions. 

In  a  similar  manner  the  properties  of  cka-boro?i  and  eka-silicon 
were  predicted,  and  the  subsequent  discovery  of  scandium  (Nilson, 
1879),  and  germanium  (Winkler,  1886),  whose  properties  were 
found  to  closely  accord  with  these  hypothetical  elements,  formed 
an  additional  demonstration  of  the  truth  of  the  periodic  law. 


The  Periodic  Classification  125 

There  are  at  present  two  elements,  however,  which  appear  not 
to  conform  strictly  to  this  periodic  classification.  These  are  the 
elements  argon  and  tellurium.  The  atomic  weight  of  argon 
according  to  most  recent  determination  is  39.92,  while  that  of 
potassium  is  39.15.  Now  the  periodic  system  requires  that  the 
atomic  weight  of  argon  shall  be  below  and  not  above  that  of 
potassium.  Again,  the  latest  determinations  of  the  atomic  weight 
of  tellurium  give  127.6,  as  against  126.85  for  iodine  ;  while  in 
order  to  conform  to  the  periodic  system  the  atomic  weight  of 
tellurium  should  be  below  that  of  iodine.  Whether  these  two 
cases  will  prove  to  be  true  exceptions,  or  whether  future  investiga- 
tions will  show  that  the  atomic  weights  here  given  are  not  the 
true  ones,  time  alone  will  show.  It  must  be  borne  in  mind,  how- 
ever, that  both  argon  and  tellurium  are  elements  which  it  is 
extremely  difficult  to  obtain  in  a  state  of  absolute  purity,  and  there 
is  considerable  probability  that  in  the  latter  case  the  element  in  a 
pure  state  has  never  yet  been  obtained. 

The  position  which  should  be  given  to  hydrogen  in  the  periodic 
system  has  been  the  subject  of  much  discussion.  It  will  be 
noticed  that  in  the  table  it  is  placed  with  a  query  in  Group  I.  and 
again  in  Group  VII.  ;  its  univalent  character  suiting  either  position 
equally  well.  The  chief  argument  in  favour  of  placing  it  in 
Group  I.  is  its  electro-positive  character,  in  which  it  strongly  re- 
sembles the  elements  lithium,  sodium,  potassium,  £c.,  metals 
which  may  be  substituted  for  hydrogen  atom  for  atom ;  the 
"salts  of  hydrogen"  (i.e.  acids\  and  the  metallic  salts  resembling 
each  other  when  regarded  from  the  ionic  standpoint. 

The  arguments  in  favour  of  assigning  it  a  position  at  the  head 
of  Group  VII.  are  more  numerous,  and  may  be  briefly  summarised 
as  follows  :* — 

1.  Its  gaseous  character  and  low  boiling-point. 

2.  Absence  of  any  metallic  properties. 

3.  The  diatomic  nature  of  its  molecules  H2  (while  many  of  the  alkali  metals 
are  monatomic). 

4.  Readiness  with  which  H  is  substituted  by  Cl,  Br,  or  I,  in  organic  com- 
pounds. 

5.  If  placed  in  Group  I.  a  series  of  six  blank  spaces  is  left,  for  as  many  un- 
known elements,  whose  atomic  weights  must  all  fall  between  H  =  i  and  He  =  3. 96. 

6.  The  numerical  difference  between  H  =  i  and  F  =  ig  is  18  units,  which  is 
much  closer  to  the  average  of  about  16  units  than  that  between  H  =  i  and 
Li  =  7,  which  is  only  6  units. 

*  Masson,  Chem.  News,  vol.  Ixxiii.  p.  283. 


CHAPTER  XIII 
GENERAL   PROPERTIES   OF   LIQUIDS 

UNDER  this  head  the  following  subjects  will  be  considered  : — 

1.  The  passage  of  liquids  into  vapours  or  gases. 

2.  The  passage  of  liquids  into  solids. 

3.  Solution,. 

1.  The  Passage  of  Liquids  into  Gases.  Evaporation  and 
Boiling. — Just  as  in  the  gaseous  condition,  so  in  the  liquid  state, 
the  molecules  are  in  a  state  of  motion  :  in  the  liquid  state,  however, 
the  mean  kinetic  energy  of  the  molecules  is  unable  to  overcome  the 
force  of  their  mutual  attraction.  Some  of  the  molecules  have  a 
smaller  kinetic  energy  (that  is,  a  lower  temperature),  and  others 
a  greater  kinetic  energy,  than  the  average  ;  and  when  in  the  course 
of  their  movements  the  latter  strike  the  surface  of  the  liquid  and 
break  through  it,  they  continue  their  movements  in  the  space 
above,  as  gaseous  molecules.  If  the  space  into  which  they  wander 
be  unlimited,  that  is,  if  the  liquid  be  freely  exposed  to  the  air,  these 
molecules  escape  away  altogether,  and  consequently  the  liquid 
diminishes  in  quantity.  This  process  is  known  as  evaporation, 
and  as  the  molecules  which  so  leave  the  liquid  are  those  having 
the  highest  temperature,  it  follows  that  the  temperature  of  the 
liquid,  which  is  the  average  temperature  of  the  molecules,  will  fall. 
The  more  completely  the  molecules  that  so  escape  from  the  surface 
of  a  liquid  are  prevented  from  falling  back,  that  is,  the  more  rapidly 
they  are  swept  away  from  the  immediate  neighbourhood  of  the 
liquid,  the  more  quickly  will  this  escape  of  molecules  take  place, 
and  therefore  the  greater  will  be  the  fall  of  temperature  that  results 
from  evaporation.  Thus,  if  a  quantity  of  liquid,  say  water,  be 
exposed  in  a  dish  so  that  a  current  of  air  is  blown  across  the  sur- 
face, the  rate  of  evaporation  is  increased,  and  the  temperature  con- 
sequently falls  lower  than  if  the  water  be  merely  placed  in  a  still 

atmosphere  ;  similarly,  if  the  water  be  placed  in  a  vacuum  the  rate 

126 


Evaporation 


127 


of  evaporation  is  increased,  because  the  molecules  that  escape  from 
the  surface  of  the  liquid  are  not  impeded  in  their  motions  by 
collisions  with  the  molecules  of  air. 

This  fall  of  temperature  resulting  from  evaporation  may  be 
readily  seen  by  enveloping  the  bulb  of  a  thermometer  in  a  piece  of 
thin  muslin,  and  moistening  it  with  water.  If  such  a  thermometer 
be  placed  by  the  side  of  a  naked  thermometer,  it  will  be  seen  that 
the  mercury  will  fall  lower  in  the  one  that  is  moistened,  and  the 
difference  will  be  still  more 
marked  if  the  instruments 
are  placed  in  a  draught, 
whereby  the  evaporation  ot 
the  water  from  the  muslin 
is  accelerated. 

If  the  space  above  the 
liquid  be  limited,  molecules 
still  continue  to  escape 
from  the  surface  ;  but  a 
state  of  equilibrium  is  soon 
established,  when  as  many 
are  thrown  back  again  by 
rebounding  from  one  an- 
other and  from  the  walls 
of  the  containing  vessel  as 
leave  the  surface  in  a  given 
time.  Under  these  con- 
ditions the  enclosed  space  ^  ;> 
is  said  to  be  saturated  with 
the  vapour  of  the  liquid. 
The  number  of  molecules 
which  escape  from  the  sur- 
face depends  upon  the  tem- 
perature, and  is  independent  of  the  pressure,  for  if  the  volume 
of  a  saturated  vapour  be  forcibly  diminished,  it  merely  results 
in  the  condensation  of  a  portion  of  the  vapour  ;  and  if  ex- 
panded, a  corresponding  vaporisation  of  an  additional  quantity 
of  the  liquid,  the  pressure  remaining  always  constant.  The  num- 
ber of  molecules  that  re-enter  the  liquid  is  determined  by  the 
number  and  the  velocity  of  those  that  exist  as  gaseous  molecules 
in  a  unit  volume.  But  the  pressure  exerted  by  a  gas  is  caused  by 
the  number  and  velocity  of  the  molecules  in  a  given  volume,  hence 


FIG.  13. 


128  Introductory  Outlines 

the  condition  of  equilibrium  is  set  up,  when  the  vapour  above  the 
liquid  exerts  a  definite  pressure,  which  pressure  will  be  constant 
for  any  given  temperature.  The  pressure  exerted  by  a  vapour 
under  these  conditions  is  termed  the  'vapour-tension  of  the  liquid. 
The  fact  that  the  vapour  given  off  from  a  liquid  exerts  pressure 
may  readily  be  experimentally  illustrated  by  means  of  the  apparatus 
seen  in  Fig.  13.  Three  glass  tubes,  A,  B,  and  C,  about  one  metre 
long,  are  completely  filled  with  mercury  and  inverted  in  a  trough 
of  the  same  liquid.  The  mercury  will  sink  to  the  same  level  in 
each  tube,  the  length  of  the  mercury  column  representing  the 
atmospheric  pressure  at  the  time.  Into  two  of  these  barometer 
tubes,  B  and  C,  a  few  drops  of  water  are  introduced,  when  it  will  be 
found  that  the  mercury  is  depressed,  as  indicated  in  B,  below  the 
level  at  which  it  previously  stood.  This  depression  of  the  mercury 
column  represents  the  tension  of  the  vapour  of  the  water  for  the 
particular  temperature  at  which  the  experiment  is  made.  If  tube 
C  be  surrounded  by  a  wider  glass  tube,  through  which  steam  from 
a  small  boiler  is  passed,  it  will  be  noticed  that  as  the  temperature 
of  the  water  in  the  tube  rises,  the  mercury  is  more  and  more  de- 
pressed, thus  showing  that  the  tension  of  the  vapour  increases  with 
rise  of  temperature.  As  soon  as  the  steam  circulates  freely  and  is 
escaping  at  the  bottom  of  the  wide  tube,  in  other  words,  as  soon 
as  the  temperature  of  the  enclosed  water  in  tube  C  reaches  100°, 
i.e.  the  temperature  of  the  steam  surrounding  it,  the  mercury  in 
the  tube  will  be  depressed  to  the  level  of  that  in  the  trough.  The 
tension  of  the  vapour  within  the  tube,  under  these  circumstances, 
is  therefore  equal  to  the  atmospheric  pressure. 

If,  instead  of  introducing  water  into  the  barometer  tube,  ether 
were  employed,  and  a  stream  of  vapour  from  boiling  ether  were 
passed  through  the  outer  tube,  it  would  be  seen  that  when  the  ether 
within  the  tube  reached  the  temperature  of  the  vapour  from  the 
boiling  ether,  namely,  35°,  the  mercury  would  again  be  depressed 
to  the  level  of  that  in  the  trough  ;  that  is,  the  tension  of  the  ether 
vapour  would  then  be  equal  to  the  pressure  of  the  atmosphere.  We 
see,  therefore,  that  when  water  is  heated  to  its  boiling-point,  viz., 
1 00°,  the  tension  of  its  vapour  is  equal  to  the  atmospheric  pressure  ; 
and  when  ether  is  heated  to  its  boiling-point,  viz.,  35°,  the  pressure 
exerted  by  its  vapour  is  equal  to  the  pressure  of  the  atmosphere. 
The  boiling-point  of  a  liquid  may  therefore  be  defined  as  the 
temperature  at  which  the  vapour-pressure  is  equal  to  the  pressure 
of  the  atmosphere.  As  soon  as  this  point  is  passed,  the  kinetic 


Boiling- Points  of  Liquids  1 29 

energy  of  the  molecules  has  been  so  much  augmented  by  the 
supply  of  external  heat,  that  it  is  able  to  overcome  the  force  of 
their  mutual  attractions,  and,  consequently,  the  molecules  freely 
pass  away  from  the  surface  of  the  liquid. 

As  will  be  seen  from  the  illustrations  given,  namely,  water  and 
ether,  the  temperatures  at  which  the  vapours  of  different  liquids 
exert  a  pressure  equal  to  that  of  the  atmosphere  are  widely  different. 
This  fact  will  be  still  more  evident  from  the  following  table,  giving 
the  temperatures  at  which  the  vapour  pressure  of  various  liquids  is 
equal  to  the  standard  atmospheric  pressure  : — 

Liquid  hydrogen       ....  ~253° 

Liquid  oxygen —182.5° 

Liquid  nitrous  oxide          ...  —   89.8° 

Liquid  sulphur  dioxide     ...  —    10° 

Ethyl  chloride +11° 

Carbon  disulphide    ....  47° 

Water 100° 

Aniline 182° 

Mercury 358° 

Since  the  boiling-point  of  a  liquid  is  that  temperature  at  which 
its  vapour-tension  is  equal  to  the  atmospheric  pressure,  it  will  be 
evident  that,  if  the  latter  increases  or  decreases,  the  temperature 
necessary  to  produce  an  equal  vapour-pressure  must  also  rise  or 
fall ;  in  other  words,  the  boiling-point  of  a  liquid  is  dependent  upon 
the  pressure.  If  a  quantity  of  water,  no  warmer  than  the  hand,  be 
placed  beneath  the  receiver  of  an  air-pump,  which  is  then  quickly 
exhausted,  the  water  will  be  seen  to  enter  into  violent  ebullition. 
It  does  this  when  the  pressure  within  the  receiver  is  reduced  to 
the  point  at  which  it  is  equal  to  the  tension  of  aqueous  vapour  at 
the  temperature  taken. 

For  this  reason  water  boils  at  a  lower  temperature  in  high 
altitudes  than  at  the  sea-level ;  and  as  the  vapour-tension  of  water 
at  various  temperatures  has  been  experimentally  determined,  we 
can,  by  ascertaining  the  boiling-point  of  water  at  any  particular 
altitude,  calculate  the  atmospheric  pressure,  and  consequently  the 
height  above  the  sea-level. 

Many  liquids  when  heated,  especially  in  glass  vessels  that  have 
been  carefully  cleansed,  may  be  raised  several  degrees  above  the 
boiling-point  without  ebullition  taking  place.  The  liquid  under 
these  circumstances  assumes  a  pulsating  movement,  which  con- 

I 


130  Introductory  Outlines 

tinues  for  a  short  time,  when  a  burst  of  vapour  is  suddenly  evolved 
with  violence,  and  the  temperature  at  once  drops  to  the  boiling- 
point.  The  liquid  then  becomes  quiescent,  and  again  as  the 
temperature  rises  the  pulsating  movement  begins,  ending  once 
more  in  an  explosive  evolution  of  vapour.  This  succussive  boiling, 
or  bumping,  is  sometimes  sufficiently  violent  to  cause  the  fracture  of 
the  vessel.  In  order  to  experimentally  ascertain  the  boiling-point 
of  a  liquid,  the  thermometer,  for  this  reason,  is  not  immersed  in 
the  liquid,  but  is  suspended  in  the  vapour,  the  temperature  of 
which  remains  constant  throughout  these  irregularities  in  the 
boiling. 

Latent  Heat  Of  Vaporisation. — When  a  liquid  is  heated,  its 
temperature  rises,  as  indicated  by  the  thermometer,  until  a  certain 
point  is  reached  (the  boiling-point  of  the  liquid),  when  the  con- 
tinued application  of  heat  causes  no  further  rise  of  temperature. 
Thermometers  placed  in  the  liquid,  and  in  the  vapour,  indicate  the 
same  temperature  and  remain  constant,  and  all  further  applica- 
tion of  heat  is  unappreciated  by  these  instruments,  and  disappears 
in  changing  the  liquid  into  vapour.  The  heat  which  in  this  way  is 
absorbed  during  the  vaporisation  of  a  liquid  is  spoken  of  as  the 
latent  heat  of  vaporisation  ;  and  the  same  amount  of  heat  which 
thus  disappears  during  the  conversion  of  a  liquid  into  a  vapour 
is  again  rendered  sensible  when  the  vapour  passes  back  into  the 
liquid  state. 

The  heat  which  is  thus  said  to  become  latent  is  in  reality  con- 
verted into  kinetic  energy  ;  it  is  expended  in  imparting  to  the 
molecules  the  kinetic  energy  necessary  to  overcome  the  attractive 
forces  operating  between  them  while  in  the 
liquid  state  ;  in  other  words,  it  is  doing  the 
work  of  overcoming  cohesion  (internal  work), 
and  also  the  external  pressure  on  the  vapour 
(external  work). 

In  order  that  a  liquid  may  pass  into  a  vapour 
it  is  necessary  that  heat  be  absorbed.      We 
have  seen  (page  126)  that  a  liquid  undergoing 
spontaneous  evaporation  becomes  colder  (that 
FIG.  14.  is>  heat  is  absorbed  by  the  molecules  that  are 

converted  into  the  gaseous  state),  and  also 
that  the  more  rapidly  the  liquid  can  be  made  to  pass  into  the 
vaporous  condition,  without  supplying  external  heat,  the  lower 
will  its  temperature  fall.  Upon  this  fact  depend  a  number 


Latent  Heat  of  Vaporisation  131 

of  methods  for  the  artificial  production  of  low  degrees  of  cold. 
For  example,  ether  boils  at  35°,  but  if  a  small  quantity  of  ether 
be  placed  in  a  glass  flask  standing  upon  a  wooden  block,  upon 
which  a  few  drops  of  water  have  been  poured,  and  a  current  of  air 
from  a  bellows  be  briskly  blown  through  the  ether  (Fig.  14), 
the  temperature  of  the  ether  will  fall  so  rapidly  that  in  a  few 
moments  the  flask  will  be  frozen  to  the  block.  By  the  rapid  eva- 
poration of  liquids  with  lower  boiling-points,  the  extreme  degrees  of 
cold  necessary  for  the  liquefaction  of  such  gases  as  oxygen,  carbon 
monoxide,  air,  &c.,  are  obtained.  Thus,  liquid  methyl  chloride 


M 


FIG.  15. 

boils  at  -23°;  by  causing  it  to  rapidly  vaporise,  its  temperature 
can  be  reduced  to  -  70°.  Liquid  ethylene  in  the  same  way  falls 
to  a  temperature  of  -  120°,  and  liquid  oxygen  by  rapid  evaporation 
gives  a  temperature  as  low  as  —210°. 

The  temperature  of  water,  in  like  manner,  may  be  so  lowered  by 
its  own  rapid  evaporation,  as  to  cause  it  to  freeze.  We  have  already 
seen  that  by  reducing  the  pressure,  the  boiling-point  of  a  liquid  is 
lowered  ;  if,  therefore,  a  quantity  of  water  be  placed  in  a  vacuum, 


132 


Introductory  Outlines 


and  methods  be  adopted  to  remove  the  water  vapour  as  rapidly 
as  it  is  formed,  the  water  will  enter  into  rapid  ebullition.  The 
evaporation  will  therefore  proceed  so  rapidly,  and  consequently 
absorb  heat  so  quickly,  that  the  temperature  of  the  boiling  liquid 
will  quickly  fall  to  o°  when  it  passes  into  the  solid  state.  The 
instrument  known  as  Carre's  freezing  machine  depends  upon  this 
principle.  The  water  to  be  frozen  is  placed  in  the  glass  bottle  C 
(Fig.  15),  which  is  in  connection  with  a  metal  reservoir  R,  half 
filled  with  strong  sulphuric  acid.  This  in  its  turn  is  connected  by 
b  with  an  air-pump  P,  worked  by  the  lever 
M,  to  which  is  also  attached  a  connecting 
rod  /,  so  that  a  stirrer  within  the  reservoir  is 
kept  constantly  in  motion.  As  soon  as  the 
apparatus  is  exhausted  to  a  pressure  of  two 
or  three  millimetres,  the  water  begins  rapidly 
to  boil,  and  as  the  sulphuric  acid  absorbs  the 
water  vapour  as  rapidly  as  it  is  given  off,  the 
temperature  quickly  falls  and  the  water 
freezes. 

Fig.  16  illustrates  another  method  by  which 
the  same  result  may  be  obtained.  A  tall 
glass  vessel  is  exhausted  by  means  of  an 
ordinary  air-pump,  and  water  is  allowed 
slowly  to  enter  from  a  stoppered  funnel,  upon 
the  end  of  which  is  secured  a  short  string. 
At  the  same  time  strong  sulphuric  acid  is  ad- 
mitted by  the  second  funnel,  and  caused  to 
flow  down  a  glass  rod,  round  which  is  wound 
a  spiral  of  asbestos  thread.  The  acid  at  once 
absorbs  the  aqueous  vapour  from  the  evapo- 
rating water,  the  temperature  of  which,  there- 
fore, falls  below  the  freezing-point,  and  it 
solidifies  as  it  flows  over  the  string  into  the 
form  of  an  icicle. 

Just  as  diminution  in  pressure  lowers  the  boiling-point  of  a 
liquid,  so  increased  pressure  raises  the  boiling-point.  If  water  be 
heated  in  a  closed  iron  vessel,  as  in  a  high-pressure  steam  boiler, 
the  pressure  caused  by  its  own  vapour  raises  the  boiling-point 
many  degrees  above  100°.  There  is  a  definite  temperature,  how- 
ever, for  every  liquid,  beyond  which  the  liquid  state  is  impossible, 
whatever  may  be  the  pressure  ;  that  is  to  say,  the  liquid  when 


FIG.  16. 


Vapour- Pressures  of  Solutions  133 

heated  beyond  this  fixed  point  passes  into  the  gaseous  state,  how- 
ever great  the  pressure  may  be.  This  temperature  is  the  critical 
temperature  (see  page  79).  If  a  liquid  be  heated  in  a  sealed  and 
strong  glass  tube,  as  the  critical  temperature  is  approached  the 
surface  of  the  liquid  gradually  becomes  ill-defined,  and  finally  the 
tube  is  completely  occupied  by  transparent  vapour.  On  again 
cooling,  as  soon  as  the  critical  point  is  passed  the  contents  of 
the  tube  again  separate  into  two  distinct  layers  consisting  of  liquid 
and  gas. 

Vapour-Pressures  of  Solutions.— The  boiling-point  of  a  liquid 
is  modified  by  the  presence  in  the  liquid  of  dissolved  substances. 
If  the  substance  in  the  solution  be  less  volatile  than  the  liquid,  the 
boiling-point  is  raised.  Thus,  while  the  boiling-point  of  pure 
water  (under  the  normal  atmospheric  pressure)  is  100°,  the  tem- 
perature at  which  saturated  aqueous  solutions  of  salts  boil,  is 
considerably  higher,  thus  : — 

Containing  Grammes  of 

Water  Saturated  with  Salt  in  100  Grammes          Boiling-point, 

of  Water. 

Sodium  chloride    .  .  .  41.2  108.4° 

Potassium  nitrate .  .  .  335.1  H5-9° 

Potassium  carbonate  .  .  205.0  I33-°° 

Calcium  chloride  .  .  .  325.0  I79-5° 

The  temperature  of  the  steam  of  these  boiling  solutions,  as 
ascertained  by  suspending  a  thermometer  in  the  vapour,  appears 
to  be  the  same  as  that  from  pure  water,  as  the  thermometer  in 
all  cases  indicates  100°.  In  reality,  however,  the  temperature  is 
higher,  although  not  so  high  as  that  of  the  boiling  liquid.  The 
reason  that  the  thermometer  indicates  100°  in  all  cases  is  because 
the  water  vapour  continually  condenses  upon  the  bulb  of  the 
instrument,  covering  it  with  a  film  of  pure  water,  which  boiling 
off  from  the  bulb  indicates  only  the  boiling-point  of  the  pure 
liquid.  By  special  arrangements  this  condensation  may  be  pre- 
vented, when  it  has  been  shown  (Magnus)  that  the  temperature 
of  the  vapour,  from  such  boiling  solutions,  rises  as  the  solutions 
become  more  concentrated — that  is,  as  the  temperature  of  the 
boiling  liquids  rise.  It  has  been  already  explained  that  the  boil- 
ing-point of  a  liquid  is  that  temperature  at  which  the  vapour 
tension  is  equal  to  the  atmospheric  pressure  ;  since,  then,  the 
presence  of  dissolved  substances  raises  the  boiling-point,  it  follows 


134  Introductory  Outlines 

that  it  must  lower  the  vapour-pressure,  for  (in  the  case  of  aqueous 
solutions)  when  the  temperature  has  reached  100°  the  vapour- 
pressure  is  still  below  that  of  the  atmosphere,  for  the  liquid  does 
not  enter  into  ebullition  at  that  temperature.  Lowering  the 
vapour-pressure,  therefore,  is  synonymous  with  raising  the  boiling- 
point.  The  extent  to  which  the  vapour-pressure  of  a  liquid  is 
lowered  (or  its  boiling-point  raised)  by  dissolving  in  100  grammes 
of  it  i  gramme-molecule  of  a  given  substance  is  called  the  mole- 
cular lowering  of  the  vapour-pressure,  or  the  molecular  elevation 
of  the  boiling-point  of  that  liquid.  Now  it  has  been  found  with 
substances  which  do  not  undergo  ionic  dissociation  in  the  solvent 
employed,  and  also  which  do  not  themselves  exert  any  appreciable 
vapour-pressure  at  the  boiling-point  of  the  solvent,  that  this  mole- 
cular lowering  of  the  vapour-pressure  is  practically  a  constant. 
Thus,  for  water  the  molecular  rise  of  boiling-point  is  5.2° ;  while  for 
benzene  it  is  27.0°. 

For  example,  given  two  substances,  say  glycerol  and  sugar, 
which,  when  dissolved  in  water,  yield  solutions  which  are  non- 
electrolytes  (i.e.  these  compounds  do  not  dissociate),  and  are  also 
themselves  practically  non-volatile  at  the  boiling-point  of  water  ; 
then,  if  I  gramme-molecule  of  each  be  separately  dissolved  in 
100  grammes  of  water,  the  two  solutions  obtained  will  be  found  to 
boil  at  about  105.2°  instead  of  100°.  Or  again,  two  substances 
fulfilling  the  same  conditions  when  dissolved  in  benzene  would 
send  up  the  boiling-point  of  this  liquid  from  80.5°  to  107.5°. 

If,  on  the  other  hand,  the  substance  is  an  electrolyte — that  is, 
one  which  undergoes  ionic  dissociation  in  the  solvent,  then  the 
effect  produced  by  the  same  weight  of  substance  is  greater,  since 
the  ions  behave  as  though  they  were  molecules,  and  the  result  is 
the  same  as  though  a  larger  number  of  molecules  were  present  in 
the  solution.  Obviously  the  increase  in  the  effect  produced  will 
depend  upon  the  extent  to  which  dissociation  takes  place. 

The  following  general  laws  relating  to  the  effect  of  dissolved 
substances  upon  vapour-pressure  have  been  established  : — 

1.  The  relation  between  the  quantity  of  a  substance  in  solution 
and  the  diminution  of  the  vapour-pressure  below  that  of  the  pure 
solvent  is  the  same  at  all  temperatures. 

2.  The  diminution  of  the  vapour-pressure  of  a  liquid,  by  a  dis- 
solved substance,  is  proportional  to  the  amount  of  the  substance  in 
solution  {provided  the  substance  itself  exerts  no  appreciable  vapour- 
pressure  at  the  temperature  of  the  experiment}. 


Vapour- Pressures  of  Solutions 


135 


3.  The  molecular  lowering  of  -vapour-pressure  by  chemically 
similar  substances  is  constant;  that  is  to  say,  solutions  containing 
one  molecular  weight  in  grammes  (one  gramme-molecule}  of  such 
substances  in  equal  volumes  of  the  solvent,  give  rise  to  the  same 
diminution  of  vapour-pressure. 


FIG.  17. 

4.  The  relative  lowering  of  vapour-pressure  is  proportional  to 
the  ratio  of  the  number  of  molecules  of  the  dissolved  substance,  to 
the  total  mimber  of  molecules  in  the  sohition,  i.e.  the  sum  of  the 
number  of  molecules  of  the  dissolved  substance  and  of  the  solvent* 

*  Except  in  the  case  of  electrolytes.     See  page  109. 


136  Introductory  Outlines 

Upon  these  considerations  it  becomes  possible,  by  means  of  the 
lowering  of  the  vapour-pressure,  to  determine  the  molecular  weight 
t>f  a  substance  that  is  capable  of  being  dissolved  in  a  volatile 
liquid. 

The  apparatus  in  which  such  a  determination  is  made  is  shown 
in  dissected  form  in  Fig.  17.  A  weighed  quantity  of  the  solvent  to 
be  employed  is  contained  in  the  tube  A  which  is  inserted  in  the 
vessel  B,  which  in  its  turn  is  placed  upon  the  asbestos  support  D, 
and  heated  from  below  by  means  of  small  flames.  As  the  liquid 
in  A  boils,  its  vapour  is  condensed  by  the  condenser  indicated  at 
C1?  and  thereby  returned  to  the  vessel.  The  outer  vessel  B  also 
contains  a  small  quantity  of  the  same  liquid  which  boils  simulta- 
neously, so  that  the  inner  tube  is  thus  surrounded  by  a  jacket  filled 
with  the  hot  vapour  of  the  same  liquid  as  is  boiling  inside.  The 
vapour  from  the  boiling  liquid  in  this  jacket  vessel  is  condensed  by 
the  condenser  at  C2  and  constantly  returned.  By  means  of  a 
thermometer  the  exact  temperature  at  which  the  liquid  boils  is 
thus  ascertained,  after  which  a  weighed  quantity  of  the  substance 
whose  molecular  weight  is  to  be  determined  is  introduced  and  the 
boiling-point  again  ascertained. 

The  result  is  calculated  by  the  formula  — 


When  C  =  Constant—  namely,  the  molecular  elevation  of  the  boil- 

ing-point of  the  solvent  used  ; 
g  =  The  percentage  strength  of  the  solution  ;  and 
R  =  The  observed  rise  of  boiling-point. 

The  Passage  of  Liquids  into  Solids.—  Most  liquids,  when 
cooled  to  some  specific  temperature,  pass  into  the  solid  state  ;  the 
temperature  at  which  this  change  takes  place  is  termed  the  solidi- 
fying point.  Generally  speaking,  the  temperature  at  which  a  liquid 
solidifies  is  the  same  as  that  at  which  the  solid  again  melts  ;  but  as 
the  solidification  of  a  liquid  is  subject  to  disturbances  from  causes 
that  do  not  affect  the  melting-point,  this  is  not  always  the  case. 
Thus,  water  may  be  cooled  many  degrees  below  o°  if  it  be  pre- 
viously freed  from  dissolved  air,  and  be  kept  perfectly  still.  This 
super-  cooling  of  water  may  readily  be  illustrated  by  means  of  the 
apparatus  represented  in  Fig.  18.  This  consists  of  a  thermometer 
whose  bulb  is  enclosed  in  a  larger  bulb  containing  water,  which 
before  the  bulb  is  sealed  at  «,  is  briskly  boiled  to  expel  all  the  air. 


Solidifying  Points  of  Liquids  137 

When  the  instrument  is  immersed  in  a  freezing  mixture  the  tem- 
perature of  the  water  may  be  lowered  to  -15°  without  congela- 
tion taking  place,  but  on  the  slightest  agitation  it  at  once  solidifies 
and  the  temperature  rises  to  o°.  It  is  on  account  of  this  property 
of  water  to  suspend  its  solidification,  that  in  determining 
the  lower  fixed  point  of  a  thermometer,  the  temperature 
of  melting  ice,  and  not  that  of  freezing  water,  is  made 
use  of. 

Many  other  liquids  exhibit  suspended  solidification  to 
a  very  high  degree  ;  thus  glycerine  may  be  cooled  to  —  30° 
or  —40°  without  solidifying,  but  if  a  crystal  of  solid  gly- 
cerine be  placed  in  the  liquid  the  entire  mass  freezes,  and 
does  not  again  melt  until  a  temperature  of  15.5°  is 


reached. 


20 


Change  of  Volume  on  Solidification.— Most  liquids, 
in  the  act  of  solidifying,  contract  ;  that  is  to  say,  the  solid 
occupies  a  smaller  volume, than  the  liquid.  Consequently 
the  solid  is  specifically  denser,  and  sinks  in  the  liquid. 
Thus  100  volumes  of  liquid  phosphorus  at  44°  (the  melting- 
point)  when  solidified  occupy  only  96.7  volumes.  Water 
expands  upon  solidification,  hence  ice  is  relatively  lighter 
than  water,  and  floats  upon  the  liquid.  The  reverse 
change  of  volume  accompanies  the  change  of  state  in  the 
opposite  direction. 

Effect  of  Pressure  upon  the  Solidifying  Point  of 
Liquids. — In  the  case  of  liquids  that  contract  upon  soli- 
dification, increased  pressure  raises  the  point  of  solidifi- 
cation, and  consequently  raises  the  melting-point  of  the 
solid.  The  effect,  however,  is  extremely  small  :  thus  the 
solidifying-point  (and  melting-point)  of  spermaceti  under 
the  standard  atmospheric  pressure  is  47-7°,  while  under  p  g 
a  pressure  of  1 56  atmospheres  it  is  raised  to  50.9°. 

With  liquids  that  expand  on  solidification,  increased  pressure  has 
the  opposite  effect,  and  lowers  the  solidifying  point.  Thus,  water 
under  great  pressure  may  be  cooled  below  o°  and  still  remain  liquid  ; 
and  in  the  same  way  ice  may  be  liquefied  by  increased  pressure 
without  altering  its  temperature.  In  the  case  of  water  it  has  been 
found  that  an  increased  pressure  of  n  atmospheres  lowers  the  soli- 
difying point  by  0.007472°  ;  hence  under  a  pressure  of  135  atmos- 
pheres, the  freezing-point  of  water  (and  the  melting-point  of  ice)  is 
lowered  i°.  This  lowering  of  the  melting-point  of  ice  under  pres- 


138 


Introductory  Outlines 


sure  may  be  illustrated  by  the  experiment  represented  in  Fig.  19. 
Over  a  block  of  ice  is  slung  a  fine  steel  wire,  to  which  are  hung  a 
number  of  weights.  The  pressure  thus  exerted  upon  the  ice,  by 
lowering  the  melting-point,  causes  the  ice  to  liquefy  immediately 
beneath  the  wire,  which  therefore  gradually  cuts  its  way  through 
the  block.  But  as  the  wire  passes  through  the  mass,  each  layer  of 
water  behind  it  again  resolidifies,  being  no  longer  subject  to  the 
increased  pressure  ;  hence,  although  the  wire  cuts  its  way  com- 
pletely through  the  ice,  the  block  still  remains  intact. 
Latent  Heat  Of  Fusion. — When  a  liquid,  at  a  temperature 
above  its  solidifying  point,  is  cooled, 
a  thermometer  placed  in  the  liquid 
indicates  its  loss  of  heat  until  solidi- 
fication begins.  At  this  point  the 
temperature  remains  constant  until 
solidification  is  complete,  when  the 
thermometer  again  begins  to  fall. 
And  again,  when  a  solid,  at  a  tem- 
perature below  its  melting-point,  is 
heated,  its  temperature  rises  until 
the  melting  begins,  but  no  further 
rise  of  temperature  takes  place  by  the 
application  of  heat  until  liquefaction 
is  complete.  The  sensible  heat  that 
so  disappears  during  fusion  is  spoken 
of  as  the  latent  heat  of  fusion.  Just 
as  in  the  passage  of  liquids  into 
gases,  this  so-called  latent  heat  re- 
presents heat  that  has  ceased  to  be 
heat)  but  which  is  converted  into 
kinetic  energy  that  is  taken  up  by  the 
molecules  :  when  the  liquid  passes 
back  into  the  solid  state,  this  energy  is  again  transformed  into 
sensible  heat. 

The  fact  that  heat  is  thus  changed  into  energy,  and  so  rendered 
insensible  to  the  thermometer,  may  be  seen  by  adding  boiling  water 
to  powdered  ice.  A  thermometer  placed  in  ice  indicates  the  tem- 
perature o°,  and  although  boiling  water  is  poured  upon  it,  so  long 
as  any  ice  remains  unmelted  no  rise  of  temperature  of  the  mixture 
results,  the  heat  contained  in  the  boiling  water  being  expended  in 
doing  the  work  of  liquefying  the  ice,  and  converting  it  into  water  at  o°. 


FIG.  19. 


Solidifying  Points  of  Liquids  139 

When  such  an  experiment  is  made  more  exactly,  it  is  found  that 
i  kilogramme  of  water  at  80.25°,  when  mixed  with  I  kilogramme,  of 
ice  at  o°,  gives  2  kilogrammes  of  water  at  o°.  That  is  to  say,  the 
amount  of  heat  contained  in  a  kilogramme  of  water  at  80.25°  is 
exactly  capable  of  transforming  an  equal  weight  of  ice  at  o°  into 
water  at  o°. 

As  the  heat  required  to  raise  the  temperature  of  i  kilogramme  of 
water  from  o°  to  i°  is  the  unit  of  heat,  or  major  calorie,  we  say  that 
the  latent  heat  effusion  of  ice  is  80.25  thermal  units  or  calories. 

During  the  solidification  of  a  liquid,  the  latent  heat  of  fusion  is 
again  given  out.  The  solidification,  therefore,  only  takes  place 
gradually,  for  the  heat  evolved  by  the  congelation  of  one  portion 
is  taken  up  by  the  neighbouring  particles,  whose  solidification  is 
thereby  retarded  until  this  heat  is  dissipated.  In  the  case  of  super- 
cooled liquids  and  super-saturated  saline  solutions,  the  solidifica- 
tion takes  place  more  suddenly,  and  the  evolution  of  the  latent  heat 
is  therefore  manifest  by  a  rise  of  temperature. 

Effect  of  Substances  in  Solution  upon  the  Solidifying  Point 
Of  a  Liquid. — It  has  long  been  known  that  a  lower  degree  of  cold 
is  necessary  to  freeze  salt  water  than  fresh  ;  and  also  that  the  water 
obtained  by  remelting  ice  from  frozen  sea-water  is  so  little  salt  as 
to  be  drinkable.  Quantitative  experiments  show  that  water  con- 
taining i  per  cent,  of  common  salt  requires  to  be  cooled  to  -0.6° 
before  the  water  begins  to  freeze  ;  and,  moreover,  that  when  such 
a  dilute  solution  begins  to  freeze,  the  solid  which  separates  out  is 
not  the  salt,  but  is  pure  ice.  This  also  holds  in  the  case  of  all 
other  solvents  that  are  capable  of  being  solidified,  the  pure  solidi- 
fied solvent  alone  separating  when  the  solution  is  frozen.  For 
instance,  benzene  freezes  at  6° ;  but  if  a  small  quantity  of  any  sub- 
stance which  it  is  capable  of  dissolving  be  added  (either  a  solid  or 
liquid  substance),  it  will  be  found  necessary  to  cool  the  liquid  beloiu 
6°  before  the  benzene  begins  to  freeze.  The  effect  of  dissolved 
substances  in  lowering  the  solidifying  point  of  the  solvent  was  first 
discovered  by  Blagden  (1788),  who  formulated  the  law  that  the 
depression  of  the  freezing-point  of  aqueous  solutions  of  the  same 
substance  was  proportional  to  the  strength  of  the  solution.  By 
referring  the  lowering  of  the  solidifying  point  to  quantities  of  the 
dissolved  substances  that  were  in  molecular  proportions,  instead 
of  to  equal  weights,  it  has  been  found  that  in  the  case  of  certain 
chemically  allied  substances  the  following  general  law  holds  good : 
Solutions  containing  in  equal  volumes  of  the  solvent  quantities  oj 


140  Introductory  Outlines 

dissolved  substances  proportional  to  their  molecular  weights  have 
the  same  point  of  solidification. 

Thus,  centi-normal  solutions  of  sodium  chloride  and  potassium 
chloride  (i.e.  solutions  containing  0.585  gramme  NaCl  and  0.746 
gramme  KC1  respectively  in  one  litre  of  water)  will  begin  to  freeze  at 
the  same  fraction  of  a  degree  below  o°.  In  other  words,  the  depres- 
sion of  the  freezing-point  of  the  solvent  is  a  function  of  the  number 
of  molecules  of  the  dissolved  substance,  irrespective  of  the  nature 
of  the  molecules.  The  extent  to  which  the  freezing-point  of  a 
liquid  would  be  depressed*  by  dissolving  in  100  grammes  of  it  one 
gramme-molecule  of  any  substance  is  called  the  molecular  depres- 
sion of  the  freezing-point  of  that  liquid,  and  it  is  found  that  in  the 
case  of  all  substances  which  are  non-electrolytes,  i.e.  which  do  not 
undergo  ionisation,  this  molecular  depression  for  a  given  liquid  is 
practically  a  constant.  Thus  in  the  case  of  water,  when  the  sub- 
stance dissolved  is  a  non-electrolyte,  the  molecular  depression  is 
about  18.5°. 

In  the  case  of  substances  which  dissociate  into  their  ions  in  the 
solution,  the  molecular  depression  will  be  greater,  depending  upon 
the  degree  of  ionisation.  Thus  in  the  case  of  strong  acids,  bases, 
and  salts,  that  is,  "  electrolytes  "  which  undergo  dissociation  to  the 
highest  degree,  it  is  found  that  the  molecular  depression  is  practi- 
cally double  that  given  by  non-electrolytes.  The  ions  in  the  liquid 
acting  as  independent  molecules,  it  will  be  obvious  that  if  dissocia- 
tion is  complete  there  will  be  twice  as  many  ions  as  there  were 
molecules  of  the  compound,  and  therefore  the  effect  produced  in 
respect  of  lowering  the  freezing-point  should  be  twice  as  great. 
The  relations  thus  established  between  the  molecular  weight  of  a 
compound  and  its  influence  in  lowering  the  freezing-point  of  a 
solvent  form  the  basis  of  a  method  for  the  determination  of  mole- 
cular weights  (Raoult's  method). 

The  process  is  carried  out  in  a  tube  quite  similar  to  tube  A,  Fig. 
17  (the  side  tube  in  this  case  being  merely  closed  with  a  cork). 
A  weighed  quantity  of  the  solvent  is  introduced  into  this  tube, 
which  is  then  carefully  cooled  in  a  freezing-mixture,  the  liquid 
being  gently  stirred  by  means  of  a  wire  passing  through  a  hole  in 
the  top  cork.  The  temperature  at  which  freezing  begins  to  take 
place  is  noted.  The  tube  is  then  withdrawn  from  the  freezing- 

*  In  actually  determining  depressions  of  freezing-point,  solutions  so  strong 
as  this  cannot  be  used.  The  determination  is  made  with  dilute  solutions,  and 
the  molecular  depression  obtained  by  calculation. 


Solidifying  Points  of  Liquids  141 

mixture  and  the  solidified  portion  allowed  to  melt,  when  a  weighed 
quantity  of  the  substance  whose  molecular  weight  is  to  be  deter- 
mined is  introduced,  and  the  operation  repeated.  The  molecular 
depression  is  calculated  from  the  formula  — 


where  C=  constant  —  the  molecular  depression  of  the  freezing-point; 
g=  grammes  of  substances  in  100  grammes  of  the  solvent  ; 
and       /=the  observed  depression  of  the  freezing-point. 


CHAPTER   XIV 
SOLUTION 

A  SOLUTION  may  be  defined  as  a  homogeneous  mixture  of  either  a 
gas,  a  liquid,  or  a  solid  with  a  liquid,  this  liquid  being  termed  the 
solvent* 

Substances  that  are  capable  of  forming  such  homogeneous  mix- 
tures with  a  solvent  are  said  to  be  soluble  in  that  liquid.  The 
solution  of  matter  in  its  three  states  will  be  treated  separately. 

1.  Solution  of  Gases  in  Liquids.— When  a  gas  is  dissolved 
by  a  liquid,  the  liquid  is  said  to  absorb  the  gas,  and  although  it  is 
held  that  most  liquids  are  capable  of  absorbing  most  gases  to  a 
greater  or  less  degree,  most  of  the  investigations  in  this  direction 
have  been  made  with  the  two  liquids,  water  and  alcohol,  by  Bunsen. 

The  quantity  of  a  gas  which  a  liquid  is  capable  of  absorbing 
depends  upon  four  factors — (i)  the  specific  nature  of  the  liquid  ; 
(2)  the  nature  of  the  gas  ;  (3)  the  temperature  of  the  liquid ;  (4) 
the  pressure. 

(i.)  The  influence  of  the  solvent  may  be  seen  by  a  comparison 
of  the  quantities  of  the  same  gas  which  equal  volumes  of  water  and 
of  alcohol  are  capable  of  dissolving,  thus — 

ioo  volumes  of  water  at  o°  dissolve  179.6  volumes  of  carbon  dioxide, 
while  loo          ,,          alcohol          ,,          432.9  ,,  ,, 

(2.)  The  various  quantities  of  different  gases  which  the  same 
liquid  will  absorb  are  found  to  extend  over  a  very  wide  range, 
thus — 

ioo  volumes  of  water  at  o°  dissolve  4.114  volumes  of  oxygen, 

while  ioo  ,,  ,,  ,,  114800.0  ,,  ammonia. 

*  Mixtures  of  gases  are  sometimes  regarded  as  solutions,  one  gas  being  said 
to  be  dissolved  in  the  other.  Gases  also  are  sometimes  spoken  of  as  dissolving 
liquids  and  solids,  when  liquid  and  solid  substances  directly  vaporise  into 
them. 


Henry's  Law 


143 


(3.)  The  volume  of  any  gas  which  a  liquid  can  absorb  diminishes 
with  a  rise  of  temperature.*  This  will  be  seen  from  the  following 
table,  where  the  volumes  of  different  gases  are  given  which  100 
volumes  of  water  will  absorb  at  various  temperatures. 


Temperature. 

Carbon  Dioxide. 

Nitrous  Oxide. 

Oxygen. 

Nitrogen. 

0 

179.6 

'30-5 

4.11 

2.03 

5 

144.9 

109.3 

3-62 

1.79 

10 

118.4 

9L9 

3-25 

1.  60 

20 

90.1 

67.0 

2.83 

1.40 

It  was  at  one  time  believed  that  the  solvent  power  of  water  for 
hydrogen  was  the  same  at  all  temperatures  between  o°  and  25°. 
Recent  experiments  have  shown,  however,  that  there  is  no  excep- 
tion to  the  general  law  in  this  case  ;  thus  it  has  been  found  that 
loo  volumes  of  water — 

At    o°  dissolve  2. 1 5  volumes  of  hydrogen. 

At    5°        „        2.06 

At  10°        „        1.98         „  „ 

At  20°        „        1.84        „  „ 

When  a  solution  of  a  gas  in  water  is  heated,  the  gas  being  less 
soluble  at  the  higher  temperature  is  expelled,  and  in  most  cases 
the  whole  of  the  gas  is  driven  off  at  the  boiling  temperature. 
This,  however,  is  not  invariably  the  case  ;  for  example,  the  solution 
of  hydrochloric  acid  in  water,  when  boiled,  will  distil,  without  further 
evolution  of  gas,  when  a  solution  of  definite  strength  is  reached 
(see  Hydrochloric  Acid). 

(4.)  The  influence  of  pressure  upon  the  volume  of  a  given  gas 
which  a  liquid  can  absorb  was  discovered  by  Henry  (1803),  and  is 
known  as  Henry's  law,  namely,  The  volume  of  the  gas  absorbed  by 
a  liquid  is  directly  proportional  to  the  pressure  of  the  gas.  If  the 
pressure  be  doubled,  the  same  volume  of  liquid  will  dissolve  twice 
the  volume  of  the  gas,  the  volume  in  each  case  being  measured 
at  o°  and  760  mm.  But  since,  according  to  Boyle's  law,  the 
volume  of  a  gas  is  inversely  as  the  pressure,  this  law  may  be  thus 
stated  :  A  given  volume  of  a  liquid  will  absorb  the  same  volume  of 
a  gas  at  all  pressures. 

*  Helium,  between  certain  limits  of  temperature,  is  an  exception. 


144  Introductory  Outlines 

Thus,  if  100  volumes  of  water  at  o°  dissolve  2.03  volumes  of 
nitrogen,  under  the  standard  atmospheric  pressure  (the  volume  of 
the  gas  being  measured  at  o°  and  760  mm.),  under  twice  this 
pressure,  i.e.  two  atmospheres,  the  same  volume  will  absorb  twice 
the  volume  of  nitrogen,  viz.,  4.06  volumes  measured  at  o°  and  760 
mm.  But  4.06  volumes  of  gas  measured  at  o°  and  760  mm.  occupy 
2.03  volumes  under  a  pressure  of  two  atmospheres,  therefore  the 
liquid  dissolves  the  same  volume  of  compressed  gas  as  of  gas 
under  ordinary  pressure. 

Henry's  law  is  sometimes  stated  in  a  slightly  altered  form.  If 
the  quantity  of  gas  present  in  a  unit  volume  of  both  the  liquid  and 
the  space  above  it  be  called  the  concentration  of  the  gas,  then  the 
law  may  be  expressed  by  saying  that  under  all  pressures,  the  ratio 
of  the  concentrations  of  the  gas  in  the  liquid,  and  in  the  space  above 
it,  remains  constant.  This  ratio  is  termed  the  coefficient  of  solu- 
bility, or  the  "  solubility  "  of  the  gas  in  the  particular  liquid. 

The  term  coefficient  of  absorption,  first  introduced  by  Bunsen,  is 
the  volume  of  the  gas  measured  at  o°  and  760  rnm.,  which  is 
absorbed  by  I  cubic  centimetre  of  a  liquid  at  the  same  tem- 
perature and  pressure  ;  and  it  is  therefore  simply  the  volume 
representing  the  "  solubility  "  of  the  gas,  reduced  to  o°. 

The  solubility  of  gases  in  liquids  is  measured  by  agitating  a 
known  volume  of  liquid  with  a  measured  volume  of  the  gas,  under 
determinate  conditions  of  temperature  and  pressure.  The  apparatus 
employed  by  Bunsen,  and  known  as  Bunsen's  absorptiometer,  is 
shown  in  Fig.  20.  It  consists  of  a  graduated  tube  e,  into  which 
known  volumes  of  the  gas  and  liquid  are  introduced.  The  lower 
end  of  this  tube  is  furnished  with  an  iron  screw,  by  means  of 
which  it  can  be  securely  screwed  down  upon  an  indiarubber  pad, 
in  order  to  completely  close  the  tube  (seen  in  the  side  figure). 
The  tube  containing  the  gas  and  liquid  under  examination  is 
lowered  into  a  tall  cylinder  g  g,  in  the  bottom  of  which  is  a 
quantity  of  mercury.  The  cylinder  is  then  filled  with  water,  and 
the  cap  p  screwed  down.  The  thermometer  k  registers  the  tem- 
perature. The  apparatus  is  then  briskly  shaken,  in  order  that  the 
liquid  in  the  eudiometer  may  exert  its  full  solvent  action  upon  the 
gas,  and  on  slightly  unscrewing  the  tube  from  the  caoutchouc  pad, 
mercury  enters  to  take  the  place  of  the  dissolved  gas.  The  tube 
is  again  closed  and  the  shaking  repeated,  and  these  operations  are 
continued  until  no  further  absorption  results.  Finally,  the  volume 
of  gas  is  measured,  the  temperature  noted,  and  the  pressure 


FIG.  20. 


146 


Introductory  Outlines 


ascertained  by  reading  the  position  of  the  mercury  within  the  tube, 
and  deducting  the  height  of  the  column  from  b  to  the  surface  of 
the  mercury  a,  from  the  barometric  pressure  at  the  time  of  making 
the  experiment.  The  temperature  of  the  water  in  the  cylinder 
may  be  varied,  and  the  coefficient  of  absorption  at  different  tem- 
peratures can  thus  be  determined. 

Fig.  21  represents  a  more  modern  absorptiometer,  being  a  modi- 
fied form  of  Heidenhain  and  Meyer's  apparatus.  In  this  instru- 
ment the  measuring  tube  and  the  absorption  vessel  are  separate, 
and  it  admits  of  the  use  of  much  larger  volumes  of  liquid.  By 
means  of  the  three-way  cock  a,  the  gas  to  be  experimented  upon 
is  introduced  into  A  by  first  raising  and 
then  lowering  B ;  and  the  volume  is 
measured  when  the  levels  of  the  mer- 
cury in  A  and  B  are  coincident.  By 
means  of  the  three-way  cock  £,  the 
vessel  C,  of  known  capacity,  and  which 
is  connected  with  A  by  means  of  a 
flexible  metal  capillary  tube,  is  filled 
with  the  desired  liquid.  The  vessels  A 
and  C  are  then  put  into  communica- 
tion, and  by  raising  B  and  opening  the 
tap  c  a  definite  volume  of  the  liquid  is 
run  out  into  a  measuring  vessel,  which 
represents  the  volume  of  gas  that  enters. 
The  gas  and  liquid  are  then  thoroughly 
agitated,  after  which  the  gas  is  passed 
back  into  A  by  lowering  B^  and,  when 
A  and  C  are  in  communication,  open- 
ing the  tap  c  beneath  mercury.  By 
measuring  the  diminution  in  volume 

suffered  by  the  gas,  the  volume  absorbed  by  the  known  volume 
of  liquid  is  obtained.  The  measuring  tube  and  absorption  vessel 
are  kept  constant  at  any  desired  temperature  by  surrounding 
them  by  water,  or  with  vapours  at  known  temperatures. 

Solubility  Of  Mixed  Gases.— When  two  gases  are  mixed 
together,  the  pressure  exerted  by  each  is  the  same  as  would  be 
exerted  if  the  other  were  absent  and  the  entire  space  were 
occupied  by  the  same  mass  of  the  one.  Thus,  if  a  mixture  of 
two  gases  are  in  the  proportion  of  two  volumes  of  one  and  one 
volume  of  the  other,  the  pressure  exerted  by  the  one  present  in 


FIG.  21. 


The  Law  of  Partial  Pressures  147 

larger  proportion  will  be  twice  as  great  as  that  of  the  other ;  this 
pressure  is  termed  the  partial  pressure  of  the  gas  under  the 
circumstances,  and  obviously  the  total  pressure  of  the  mixture 
will  be  the  sum  of  the  partial  pressures  of  the  constituents.  As  the 
solubility  of  a  gas  in  a  liquid  is  proportional  to  the  pressure,  the 
solubility  of  the  gases  in  a  gaseous  mixture  will  be  influenced  by 
the  proportions  in  which  they  are  present  in  the  mixture.  This 
is  known  as  Dalton's  law  of  partial  pressures,  which  may  be  thus 
stated  :  The  solubility  of  a  gas  in  a  gaseous  mixture  is  proportional 
to  its  partial  pressure.  For  example,  the  atmosphere  consists  of 
a  mixture  of  oxygen  and  nitrogen,  in  the  proportion  of  four  volumes 
of  nitrogen  to  one  volume  of  oxygen  (in  round  numbers).  The 
partial  pressure  exerted  by  the  oxygen  is  therefore  only  one-fifth  of 
the  total  atmospheric  pressure,  and  consequently  the  amount  of 
oxygen  which  a  given  volume  of  a  liquid  is  capable  of  dissolving 
from  the  atmosphere  is  only  about  one-fifth  of  that  which  it  will 
absorb  from  pure  oxygen — in  other  words,  will  be  one-fifth  the 
absorption  coefficient  of  oxygen  for  that  liquid. 

The  application  of  the  law  of  partial  pressures  will  be  seen  in 
the  solvent  action  of  water  upon  the  atmosphere.  Taking  the 
coefficients  of  absorption  of  oxygen  and  nitrogen  for  water  as 
given  by  Bunsen — 

Oxygen  =  .04114  ;         Nitrogen  =  .02035, 

and  the  proportion  of  oxygen  to  nitrogen  in  the  air  as  one  to  four, 
by  volume,  we  get — 

£4114  =  .00823,  and  -02035  x  4  =  >OI628) 

for  the  number  of  cubic  centimetres  of  oxygen  and  nitrogen  which 
will  be  dissolved  from  the  atmosphere  by  I  cubic  centimetre  of 
water  at  o°. 

One  hundred  volumes  of  water,  therefore,  will  dissolve  2.451 
volumes  of  air,  of  which  .823  volume  is  oxygen  and  1.628  volumes 
is  nitrogen  ;  and  if  this  dissolved  air  be  again  expelled  from  the 
water  by  boiling,  the  air  so  obtained  will  contain  oxygen  and 
nitrogen  in  the  proportions — 

Oxygen 33.6 

Nitrogen      ...<,.     66.4 

100.0 


148  Introductory  Outlines 

If  a  mixture  of  oxygen  and  nitrogen  in  this  proportion  be  once 
more  dissolved  in  water,  since  the  percentage  of  oxygen  has  risen 
from  20  to  33.6,  and  the  partial  pressure  proportionately  increased, 
the  mixture  of  the  two  gases  that  will  be  dissolved  will  be  still 
richer  in  oxygen  ;  and  after  solution  in  water  for  the  third  time  the 
boiled-out  air  will  be  found  to  contain  as  much  as  75  per  cent, 
of  oxygen.  It  will  be  obvious  that  the  partial  pressure  which  de- 
termines the  extent  to  which  the  separate  gases  in  a  mixture  are 
dissolved  is  not  represented  by  the  proportion  in  which  the  gases 
are  present  before  solution,  but  that  in  which  they  exist  in  the 
gaseous  mixture  after  the  solvent  has  become  saturated. 

Henry's  law  does  not  hold  good  in  the  case  of  such  very  soluble 
gases  as  ammonia,  hydrochloric  acid,  &c.  These  gases  appear  to 
enter  into  a  true  chemical  union  with  the  water,  and  in  most  of 
these  cases  the  act  of  solution  is  attended  with  considerable  evolu- 
tion of  heat.  In  some  of  these  instances  the  deviation  from  the 
law  diminishes  with  rise  of  temperature  ;  thus  at  temperatures 
above  40°  the  absorption  of  sulphur  dioxide  obeys  the  law,  while 
in  the  case  of  ammonia  conformity  to  the  law  is  observed  at  100°. 

The  gases  dissolved  by  a  liquid  are  not  only  expelled  by  boiling, 
but  are  withdrawn  by  placing  the  solution  in  a  vacuum.  This,  in- 
deed, follows  from  Henry's  law,  for  if  the  solubility  is  proportional 
to  the  pressure,  and  the  pressure  is  nil,  the  amount  of  gas  dissolved 
must  also  be  nil. 

The  molecules  of  gas  dissolved  by  a  liquid  are  regarded  as  being 
held  by  some  attractive  forces  exerted  between  them  and  the  mole- 
cules of  the  liquid  ;  in  the  course  of  their  movements,  gas  molecules 
are  constantly  leaving  and  entering  the  liquid,  and  equilibrium  is 
established  when  the  same  number  enter  and  escape  from  the 
surface  of  the  liquid  in  the  same  time.  When  the  pressure  is  in- 
creased, more  gas  molecules  strike  the  surface  in  a  unit  of  time,  and 
consequently  a  greater  volume  is  absorbed.  When  a  solution  of  a 
soluble  gas  is  placed  in  an  atmosphere  of  another  gas,  the  dissolved 
gas  continues  to  leave  the  liquid  until  equilibrium  is  established 
between  the  pressure  exerted  by  the  gas  so  leaving  and  the  amount 
remaining  in  solution.  For  this  reason  a  solution  of  ammonia, 
when  left  exposed  to  the  air,  rapidly  becomes  weaker,  owing  to 
the  escape  of  the  dissolved  gas  into  the  atmosphere.  This  process 
is  accelerated  if  a  stream  of  a  less  soluble  gas  be  caused  to  bubble 
through  the  solution. 

Solubility  Of  Liquids  in  Liquids.— The  solubility  of  liquids  in 


Solution  149 

liquids  may  be  divided  into  two  orders.  First,  cases  in  which  the 
degree  of  solubility  of  one  in  the  other  is  unlimited  ;  and  second, 
cases  where  the  extent  of  the  solubility  is  limited,  or  where  the 
liquids  are  said  to  be  partially  miscible.  Two  liquids  whose 
solubility  in  each  other  is  unlimited  are  said  to  be  miscible  in  all 
proportions ;  thus  alcohol  and  water  are  capable  of  forming  a 
homogeneous  mixture  when  added  together  in  any  proportion. 

In  the  second  class,  where  the  solubility  of  two  liquids  for  each 
other  is  Irmited,  it  is  found  that  each  liquid  is  capable  of  dissolving 
some  of  the  other.  Thus,  if  equal  volumes  of  ether  and  water  are 
shaken  together,  the  liquids  will  afterwards  separate  out  into  two 
distinct  layers,  one  floating  upon  the  other.  The  heavier  layer  at 
the  bottom  is  an  aqueous  solution  of  ether,  containing  about  10  per 
cent,  of  ether  ;  while  the  upper  liquid  is  an  ethereal  solution  of  water 
containing  about  3  per  cent,  of  water.  The  presence  of  ether 
dissolved  in  the  water  may  be  proved  by  separating  the  two  layers 
and  gently  heating  the  aqueous  liquid  in  a  small  flask,  when  the 
dissolved  ether  will  be  expelled  and  can  be  inflamed.  The  pre- 
sence of  the  water  in  the  ether  is  also  readily  proved,  either  by 
introducing  into  the  liquid  a  small  quantity  of  dehydrated  copper 
sulphate,  which  will  rehydrate  itself  at  the  expense  of  the  water  in 
the  ether,  and  be  changed  from  white  to  blue  ;  or  by  placing  in  the 
ethereal  liquid  a  fragment  of  sodium,  which  decomposes  the  dis- 
solved water  with  the  liberation  of  hydrogen. 

Another  illustration  of  two  partially  miscible  liquids  is  seen  in 
the  case  of  a  strong  aqueous  solution  of  potassium  carbonate 
and  strong  ammonia,  which  is  of  special  interest  as  being  the 
only  example  at  present  known  of  two  aqueous  solutions  of  in- 
organic substances  which  exhibit  this  phenomenon.*  Thus,  when 
strong  aqueous  ammonia  (sp.  gr.  0.880)  is  added  to  a  concen- 
trated solution  of  potassium  carbonate,  the  two  liquids  separate 
from  each  other  in  two  distinct  layers,  the  upper  layer  consisting 
of  ammonia  which  has  taken  up  a  certain  amount  of  potassium 
carbonate,  while  the  lower  liquid  consists  of  a  solution  of  potassium 
carbonate  which  has  dissolved  a  definite  quantity  of  ammonia. 

In  most  cases  the  solubility  of  liquids  in  liquids  is  increased  by 
rise  of  temperature,  although  in  some  it  is  decreased.  As  an 
example  of  the  former,  the  case  of  these  two  aqueous  liquids  may 
be  quoted.  If  the  temperature  be  raised  then  the  solubility  of 
each  of  these  solutions  in  the  other  steadily  increases,  and  the 

*  Newth,  Trans.  Chem.  Soc.,  1900,  p.  775. 


150  Introductory  Outlines 

composition  of  the  two  layers  will  therefore  gradually  approximate 
until  a  point  is  reached  at  which  they  become  identical.  This 
point  is  arrived  at  when  the  temperature  reaches  about  43°,  and  at 
this  temperature,  therefore,  the  two  liquids  are  miscible  in  all 
proportions.  If  this  liquid  be  now  cooled  below  this  temperature, 
separation  into  the  two  phases ,  as  it  is  termed,  at  once  begins,  and 
the  liquid  gradually  becomes  milky  or  turbid  owing  to  the  pre- 
cipitation from  it  of  the  heavier  solution  in  minute  drops. 

An  instance  of  decreased  solubility  by  rise  of  temperature  is 
seen  in  the  case  of  a  mixture  of  triethylamine  and  water.  If  equal 
volumes  of  these  liquids  be  mixed  together,  at  a  temperature  below 
20°,  complete  solution  takes  place,  and  a  single  homogeneous 
liquid  results.  On  warming  the  solution  it  becomes  turbid,  owing 
to  the  separation  of  the  liquid  into  two  portions,  which  ultimately 
settle  out  as  two  distinct  layers.  As  the  temperature  of  the  solu- 
tion approaches  20°,  the  liquid  becomes  very  sensitive  to  a  slight 
rise  of  temperature,  the  heat  of  the  hand  being  sufficient  to  cause 
turbidity  in  the  solution. 

It  will  be  evident,  therefore,  from  these  considerations  that  the 
distinction  between  liquids  which  are  miscible  in  all  proportions 
and  those  which  are  only  partially  miscible  is  after  all  only  an 
arbitrary  one,  the  difference  being  simply  a  function  of  the  tem- 
perature. It  is,  nevertheless,  a  convenient  distinction  to  make,  so 
long  as  we  understand  that  it  refers  to  liquids  at  the  ordinary 
temperature. 

Solution  of  Solids  in  Liquids.— When  a  solid  is  immersed  in 
a  liquid,  the  forces  which  oppose  the  solution  of  the  solid  are  the 
attractive  forces  exerted  by  the  molecules  of  the  solid  upon  each 
other  and  those  of  the  liquid  upon  themselves.  The  forces  that 
tend  to  effect  solution  are  the  attractive  forces  exerted  by  the 
molecules  of  the  liquid  upon  the  molecules  of  the  solid,  and  the 
kinetic  energy  of  the  molecules. 

By  the  action  of  the  liquid,  the  attractive  force  between  the  mole- 
cules of  the  solid  is  diminished,  and  those  molecules  nearest  the 
surface,  by  their  own  energy  and  the  attraction  exerted  by  the 
liquid,  pass  into  and  through  the  liquid.  In  the  course  of  their 
movements,  these  sometimes  return  to  the  solid,  and  a  condition 
of  equilibrium  is  finally  established  when  as  many  molecules  leave 
the  surface  of  the  solid  as  return  to  it  in  a  given  time.  Under  these 
circumstances  the  solution  is  said  to  be  saturated  with  respect  to 
the  particular  solid. 


Solution  151 

Saturated  Solutions.— The  amount  of  solid  held  in  solution  by 
the  liquid  when  the  latter  is  saturated  depends  upon  the  tempera- 
ture, for  if  the  temperature  be  raised,  the  kinetic  energy  of  the 
molecules  is  increased,  and  consequently  an  increased  number  will 
become  detached  from  the  solid.  As  a  general  rule,  therefore,  the 
solubility  of  a  solid  in  a  liquid  is  increased  by  rise  of  temperature. 
A  saturated  solution  at  a  given  temperature  may  be  obtained  in 
two  ways,  namely,  by  maintaining  the  liquid  at  that  temperature 
and  stirring  into  it  an  excess  of  the  solid,  until  no  more  of  it  is  dis- 
solved ;  or  by  dissolving  a  larger  quantity  of  the  solid  at  a  higher 
temperature,  and  allowing  the  solution  to  stand  in  contact  with  an 
excess  of  undissolved  solid,  until  the  temperature  falls  to  the  specified 
point.  During  the  cooling  the  amount  of  solid  that  the  liquid  had 
taken  up,  over  and  above  that  which  was  necessary  to  saturation 
at  the  lower  temperature,  is  deposited. 

Supersaturated  Solutions. — The  condition  of  saturation  can 
only  be  determined  when  an  excess  of  the  undissolved  solid  is 
present  in  the  liquid  ;  for  when  a  solution,  which  is  not  in  contact 
with  any  of  the  undissolved  solid,  is  brought  to  the  point  of  satura- 
tion, either  by  cooling  or  by  evaporation  of  the  liquid,  it  frequently 
happens  that  no  separation  of  solid  takes  place.  Solutions  can  in 
this  way  be  obtained,  in  which  a  larger  amount  of  the  solid  remains 
dissolved  at  a  given  temperature  than  corresponds  to  the  amount 
required  to  form  a  saturated  solution  at  that  temperature  :  such 
solutions  are  said  to  be  supersaturated.  If  into  such  a  supersatu- 
rated solution  a  fragment  of  the  solid  be  introduced,  molecules  of 
the  dissolved  solid  at  once  deposit  themselves  upon  it,  and  this 
separation  of  the  dissolved  substance  continues  until  the  solution 
reaches  a  state  of  concentration  corresponding  to  its  normal  satura- 
tion at  the  particular  temperature.  The  introduction  into  a  super- 
saturated solution  of  a  particle  of  the  solid,  in  respect  to  which  the 
solution  is  supersaturated,  is  the  only  sure  method  of  bringing 
about  the  separation  of  the  excess  of  the  dissolved  substance  ;  such 
a  solution,  therefore,  may  be  preserved  for  an  indefinite  time,  if  it  be 
kept  in  an  hermetically  sealed  vessel.  Minute  particles  of  the  solid 
towards  which  a  solution  is  supersaturated,  that  might  be  present 
in  the  dust  of  the  air,  falling  into  such  a  solution,  will  determine 
the  deposition  of  the  dissolved  solid. 

The  phenomenon  of  supersaturation  is  strictly  analogous  to  that 
of  supercooling,  or  the  suspended  solidification  of  fused  solids,  and 
is  exhibited  most  readily  by  salts  containing  water  of  crystallisa- 


152 


Introductory  Outlines 


tion,  such  as  sodium  acetate,  NaC2H3O2,3H2O  ;  sodium  thiosul- 
phate,  Na2S2O3,5H2O  ;  and  sodium  sulphate,  Na2SO4,10H2O. 
Thus,  if  a  small  quantity  of  water  be  poured  into  a  flask  nearly 
filled  with  crystallised  sodium  thiosulphate  (the  so-called  "  Hypo  " 
of  the  photographer),  and  the  mixture  be  warmed  by  immersion  in 


240 


O"      2O'      2O"    SO' 


°     03°.      7O°      QO°    90°    2OQ 


Temp  erature  . 


FIG.  22. 


hot  water,  the  whole  of  the  salt  will  dissolve  ;  and  if  the  solution  be 
then  allowed  to  cool  undisturbed,  it  will  assume  the  ordinary  tem- 
perature,, and  still  remain  fluid.  If  into  the  supersaturated  solution 
a  crystal  of  the  salt  be  dropped,  the  excess  of  salt  present  in  solution 


Solution  153 

beyond  the  normal  quantity  for  saturation  at  that  temperature  will 
crystallise  out,  and  so  great  is  this  excess  that  the  contents  of  the 
flask  will  appear  practically  solid. 

The  different  solubility  of  various  solids  in  the  same  liquid,  and 
the  increased  solubility  by  rise  of  temperature,  is  graphically  shown 
in  Fig.  22,  where  the  solubility  curves  of  five  salts  in  water  are 
represented.  The  abscissas  indicate  temperatures,  and  the  ordi- 
nates  the  number  of  parts  of  salt  dissolved  by  loo  parts  of  water. 

Thus  at  o°  1 80  grammes  of  water  will  dissolve  35.7  parts  of 
sodium  chloride,  and  as  the  temperature  is  raised  the  quantity  of 
salt  which  the  water  will  dissolve  very  slowly  increases,  until  at  100° 
the  amount  is  nearly  40  parts  :  sodium  chloride  is  therefore  nearly 
equally  soluble  in  water  at  all  temperatures. 

In  the  case  of  potassium  nitrate,  100  grammes  of  water  at  o°  will 
only  dissolve  13.3  grammes  of  the  solid,  but  as  the  temperature 
rises  the  amount  capable  of  being  dissolved  by  this  quantity  of 
water  very  rapidly  increases,  until  at  75°  150  grammes  are  dissolved. 
Lead  nitrate  is  more  soluble  than  potassium  nitrate  between  o°  and 
50°,  but  above  this  point  it  is  not  so  soluble  as  the  other,  hence  the 
two  curves  intersect  at  that  temperature.  The  solubility  of  sodium 
sulphate  in  water  appears  at  first  sight  to  be  anomalous.  The 
solubility  at  first  rapidly  increases  with  rise  of  temperature  from 
o°,  and  reaches  a  maximum  at  a  point  between  33°  and  34°,  when 
it  gradually  diminishes  with  further  rise  of  temperature.  This 
behaviour  is  in  reality  due  to  the  fact  that  we  are  not  dealing  with 
one  and  the  same  substance  throughout  the  experiment.  Sodium 
sulphate  exists  as  a  solid  in  at  least  three  forms,  namely,  the 
decahydrate,  Na2SO4,10H2O  (ordinary  Glauber's  salt)  ;  the  hepta- 
hydrate,  Na2SO4,7H2O  ;  and  the  anhydrous  salt,  Na2SO4.  The 
first  portion  of  the  curve  (Fig.  23)  represents  the  solubility  of 
Glauber's  salt  ;  thus,  at  20°  such  an  amount  of  this  decahydrated 
salt  is  dissolved,  that  the  solution  contains  20  grammes  of  Na2SO4 
in  100  grammes  of  water.  The  solubility  of  this  salt  rapidly  rises 
until  34°  is  reached,  at  which  temperature  the  salt  melts,  and  is 
then  miscible  with  water  in  all  proportions.  The  melted  salt  con- 
tains 78.8  parts  of  Na2SO4  in  100  parts  of  water,  which  is  indicated 
as  the  highest  point  upon  its  curve  : — 

Na«SO4.  10H..O. 

23  +  23  +  32  +  64        (2  +  16)  x  10 

180  :  100=  142  :  78.8. 
142  1 80 


154 


Introductory  Outlines 


The  decahydrated  salt  is  unable  to  exist  as  such  at  temperatures 
higher  than  34°,  and  when  the  melted  salt  is  heated  above  this 
point  it  is  converted  into  the  anhydrous  salt  and  water  satu- 
rated with  the  salt ;  therefore  above  34°  it  is  not  possible  to  have 
a  solution  of  sodium  sulphate  in  contact  with  solid  Glauber's 
salt.  It  can,  however,  be  in  contact  with  the  anhydrous  salt,  and 
the  second  portion  of  the  curve  expresses  the  solubility  of  this  com- 
pound in  water,  which  slowly  diminishes  as  the  temperature  rises. 


^  0° 

i- 


^ 

'i* 


- 


ao 


70°      20"     SO'    -W"     SO°     6O°      7O°     80°     90"   JOO' 

Temperature,. 

FIG.  23. 

Paradoxical  as  it  may  at  first  appear,  it  is  possible,  by  gradually 
cooling  solutions  of  a  salt  in  water,  to  cause  them  to  become  either 
more  concentrated  or  more  dilute  according  to  circumstances.  It 
has  been  already  explained  (page  139)  that  when  a  dilute  solution 
of  a  salt  in  water  is  cooled  below  o°,  ice  only  separates  out. 
Obviously,  therefore,  the  solution  that  remains  is  more  con- 
centrated than  at  first,  and  its  freezing-point  will  consequently  be 
lowered.  If  the  cooling  be  continued,  more  and  more  ice  separates 
out,  and  the  remaining  liquid  becomes  gradually  more  and  more 
concentrated  until  at  length  a  point  is  reached  when  the  solution 
is  saturated  for  that  particular  temperature.  If  cooled  below  this 
point  ice  still  separates,  but  as  the  solution  would  then  be  super- 


Osmotic  Pressure  155 

saturated,  salt  also  separates  out  ;  and  the  composition  of  the 
mixture  of  ice  and  salt  which  thus  separates  is  the  same  as  that  of 
the  remaining  solution  ;  in  other  words,  the  solution  freezes  as 
though  it  were  a  pure  chemical  compound  of  the  water  and  the 
salt  in  solution.*  Such  a  solution  is  known  as  a  constant-freezing 
solution,  or  sometimes  a  cryohydric  solution,  and  is  comparable 
with  such  constant-boiling  mixtures  as  are  obtained  by  distilling 
either  nitric  or  hydrochloric  acids  (see  pages  239  and  367). 

If  now,  instead  of  starting  with  a  dilute  solution,  a  concentrated 
solution  is  gradually  cooled,  at  some  particular  temperature 
(depending  upon  the  degree  of  concentration  at  first)  the  solu- 
tion will  become  saturated  for  that  temperature.  Further  cooling 
below  this  point  will  then  cause  the  solution  to  deposit  some  of  the 
salt  ;  and  the  liquid,  although  still  a  saturated  one  as  respects  this 
lower  temperature,  will  be  more  dilute.  As  the  cooling  continues, 
the  separation  of  the  salt  continues,  and  the  solution  therefore 
becomes  more  and  more  dilute  (still  remaining  saturated  for  the 
lower  temperatures)  until  the  point  is  reached  when  the  solution  is 
of  such  a  strength  that  any  further  separation  of  salt  (i.e.  dilution) 
would  yield  a  liquid  which  is  below  its  own  freezing-point.  That 
is  to  say,  the  water  itself  now  begins  to  freeze  and  separate  along 
with  the  salt,  and  at  this  point  the  solution  has  reached  the  same 
constant-freezing  condition  as  in  the  former  case. 

Osmotic  Pressure. — When  a  dilute  solution  of  a  substance  in 
water  is  placed  in  a  vessel  closed  with  an  animal  membrane,  such 
as  bladder  (M,  Fig.  24),  and  the  whole  is  immersed  in  water  to 
such  a  depth  that  the  level  of  the  water  outside  is  coincident  with 
that  of  the  solution  within,  it  is  found  that  the  liquid  in  the  inner 
vessel  increases  in  volume,  as  seen  by  the  fact  that  it  gradually 
rises  in  the  narrow  stem  of  the  apparatus.  Water,  therefore,  from 
the  outer  vessel  must  have  passed  in  through  the  membrane,  and 
inasmuch  as  some  of  the  dissolved  substance  is  found  in  the  water 
of  the  outer  vessel,  some  of  the  solution  must  at  the  same  time 
have  made  its  escape  through  the  membrane.  After  the  liquid  has 
risen  to  a  certain  height  in  the  narrow  tube,  it  again  begins  to 
fall,  as  the  contained  solution  continues  to  penetrate  the  mem- 
brane. This  process  is  known  as  endosmose,  and  the  instrument 
described  is  called  an  endosmorneter. 

*  At  one  time,  indeed,  such  solutions  were  believed  to  contain  definite 
chemical  compounds  of  the  salt  with  water,  which  were  called  cryohydrates 
(Guthrie). 


156 


Introductory  Outlines 


Many  attempts  were  made  to  establish  general  relations  between 
the  height  to  which  the  liquid  rose  in  the  narrow  tube  and  the 
quantities  of  substance  in  the  solution,  but  it  was  found  impossible 
to  obtain  accurate  or  comparable  measurements,  for  not  only  were 
the  results  disturbed  by  the  effect  of  the 
constantly  changing  pressure  upon  the  rate 
at  which  the  dissolved  substance  escaped 
through  the  membrane,  but  different  ani- 
mal membranes  yielded  different  results. 

Semipermeable  Membranes.—  It  was 
first  discovered  by  Traube  (1867),  and 
afterwards  extended  by  Pfeffer  (1877),  that 
artificial  membranes,  or  pellicles,  could  be 
obtained,  which,  while  allowing  of  the  pass- 
age of  water  through  them  just  as  in  the 
case  of  animal  membranes,  unlike  these 
materials,  they  offered  a  perfect  barrier  to 
the  passage  of  many  substances  in  solu- 
tion in  the  water.  Such  pellicles  are 
known  as  semipermeable  membranes. 
The  material  that  has  been  found  most 
suitable  is  precipitated  copper  ferro- 
m  cyanide.  If  a  solution  of  copper  sulphate 
J?  (CuSO4)  be  brought  cautiously  in  contact 
with  a  solution  of  potassium  ferrocyanide 
(K4Fe(CN)6),  at  the  point  where  the  two 
liquids  meet,  a  film  or  pellicle  of  precipitated  copper  ferrocyanide 
(Cu2Fe(CN)6)  is  produced.  In  order  to  make  use  of  this  extremely 
fragile  membrane,  Pfeffer  devised  the  plan  of  precipitating  it  within 
the  walls  of  a  vessel  made  of  unglazed  porcelain.  A  small  clay 
cylindrical  cell,  after  thorough  cleansing,  was  filled  with  a  dilute 
solution  of  potassium  ferrocyanide,  and  immersed  in  dilute  copper 
sulphate.  As  these  solutions  entered  the  pores  of  the  clay,  and  there 
met,  a  membrane,  consisting  of  copper  ferrocyanide,  was  formed 
within  the  walls,  which  under  these  circumstances  was  sufficiently 
strong  to  withstand  a  pressure  of  five  or  six  atmospheres. 

If  such  a  cell,  furnished  with  a  semipermeable  membrane,  be 
employed  as  an  endosmometer,  and  a  dilute  solution,  say  of  sugar, 
be  placed  within  the  apparatus,  which  is  then  immersed  in  water, 
it  is  found  that  the  liquid  rises  in  the  narrow  tube  to  a  certain 
height  above  the  level  of  the  water  in  the  outside  vessel,  and 


FIG.  24. 


Osmotic  Pressure 


157 


remains  stationary.  Water  passes  through  the  membrane,  but  no 
dissolved  substance  passes  out.  At  first  more  water  penetrates 
inwards  than  passes  out,  hence  the  increased  volume  of  liquid  in 
the  cell ;  but  when  a  certain 
pressure  is  reached,  repre- 
sented by  the  height  to  which 
the  liquid  rises  in  the  narrow 
tube,  equilibrium  is  estab- 
lished, and  water  then  passes 
in  each  direction  at  equal 
rates.  The  pressure  at  which 
this  equilibrium  is  established 
is  called  the  osmotic  pressure 
of  the  solution. 

Fig.  25  shows  the  apparatus 
employed  by  Pfeffer.  z  is  the 
porous  cell,  in  the  walls 
of  which  the  semipermeable 
membrane  is  precipitated. 
Into  this  are  cemented  the 
glass  tubes  v  and  /,  the  latter 
being  attached,  in  the  manner 
indicated,  to  a  mercury  mano- 
meter, m.  When  the  cell  con- 
taining a  solution  is  immersed 
in  water,  the  increased  volume 
of  the  contained  liquid  that 
results  causes  a  compression 
of  the  air  enclosed  in  the 
upper  part  of  the  apparatus, 
which  consequently  drives  up 
the  mercury  in  the  little  mano- 
meter, which  thus  affords  a 
means  of  measuring  the  os- 
motic pressure  of  the  solution 
under  examination. 

The  following  laws  in  rela- 
tion to  osmotic  pressure  have 
been  established  : — 


FIG.  25. 


r.  Temperature   and   concentration   being   the  same,  different 
substances  when  in  solution  exert  different  pressures. 


158  Introductory  Outlines 

2.  For  one  and  the  same  substance,  at  constant  temperature,  the 
pressure  exerted  is  proportional  to  the  concentration. 

3.  The  pressure  for  a  solution  of  a  given  concentration  is  pro- 
portional to  the  absolute  temperature,*  the  volume  being   kept 
constant. 

4.  Equimolecular  quantities  of  different  substances  (i.e.  quanti- 
ties in  the  ratio  of  their  gramme-molecule  weights),  when  dissolved 
in  the  same  volume  of  solvent,  exert  equal  pressures  at  the  same 
temperature.t 

The  analogy  between  these  laws  and  those  relating  to  gaseous 
pressure  is  very  close.  Thus  the  second  statement  corresponds 
with  Boyle's  law,  when  we  consider  the  term  concentration  to 
denote  the  quantity  of  gas,  that  is,  the  number  of  molecules,  in  a 
given  space  ;  for  if  the  number  of  molecules  in  a  unit  space  be 
doubled,  the  gaseous  pressure  is  doubled,  and  if  the  number  of 
molecules  of  dissolved  substance  in  a  given  volume  of  water  be 
doubled,  the  osmotic  pressure  is  doubled. 

The  third  statement  corresponds  with  the  law  of  Charles  :  the 
volume  of  a  gas  is  proportional  to  the  absolute  temperature  ;  or,  if 
the  volume  be  maintained  constant,  the  pressure  exerted  by  a  gas 
is  proportional  to  the  absolute  temperature. 

Osmotic  pressure,  therefore,  just  as  gaseous  pressure,  increases 
with  rise  of  temperature  and  diminishes  with  fall  of  temperature. 

Again,  in  the  last  of  these  laws,  we  see  the  extension  of  Avogadro's 
hypothesis  into  the  region  of  solution.  Avogadro's  hypothesis 
states  that  equal  volumes  of  all  gases  contain  (under  similar  con- 
ditions) an  equal  number  of  molecules — that  is  to  say,  an  equal 
number  of  molecules  at  equal  temperatures  exert  the  same  pressure. 
But  an  equal  number  of  molecules  of  different  gases  represents  an 
amount  of  the  gases  in  the  ratio  of  their  molecular  weights,  hence 
Avogadro's  hypothesis  may  be  stated  :  equimolecular  quantities  of 
gases  at  the  same  temperature  exert  equal  pressures  ;  and  this 
statement,  as  we  have  seen,  is  only  true  of  molecules  which  do 
not  dissociate  when  they  pass  into  the  gaseous  state. 

This  close  analogy  between  the  gaseous  laws  and  those  regulat- 
ing the  behaviour  of  substances  in  dilute  solution  is  explained  on 

*  By  absolute  temperature  is  meant  the  number  of  degrees  above  -2730>C. 

t  This  is  only  true  of  those  substances  whose  molecules  neither  dissociate 
into  simpler  forms  (i.e.  non-electrolytes),  nor  associate  into  more  complex  groups 
when  in  solution. 


Osmotic  Pressure 


159 


the  assumption  that  the  molecules  of  the  dissolved  body  in  a  dilute 
solution  are  so  far  apart  that  their  mutual  attractive  forces  are 
reduced  to  a  minimum,  just  as  they  are  in  the  case  of  gaseous 
molecules,  and  that  only  such  properties  are  exhibited  by  them 
as  depend  upon  their  number  in  a  unit  space.  Further,  it  has 
been  shown  in  the  case  of  a  dilute  solution  of  sugar  that  the 
osmotic  pressure  (experimentally  determined)  is  the  same  as  the 
gaseous  pressure  that  would  be  exerted  by  the  weight  of  sugar 
present  in  the  solution,  if  it  were  converted  into  gas  and  made  to 
occupy  the  same  volume  as  that  occupied  by  the  solution  at  the 
same  temperature  ;  hence  the  general  statement 
that  the  pressure  exerted  by  a  substance  in  dilute . 
solution  (its  osmotic  pressure)  is  the  same  as 
would  be  exerted  by  the  same  amount  of  the  sub- 
stance if  it  existed  as  gas  and  occupied  the  same 
•volume  at  the  same  temperature. 

Diffusion  of  Dissolved  Substances.— If  a 
quantity  of  a  soluble  solid  substance  be  placed 
at  the  bottom  of  a  vessel,  which  is  then  filled 
with  water,  the  solid  dissolves,  and  a  layer  of  a 
strong  solution  is  formed  at  the  bottom.  In 
time,  however,  the  dissolved  substance  gradually 
diffuses  throughout  the  liquid.  This  process  of 
diffusion  may  be  illustrated  by  means  of  the 
experiment  represented  in  Fig.  26.  At  the 
bottom  of  the  tall  cylinder  is  placed  a  layer  of 
a  strong  solution  of  ferric  chloride,  and  upon 
this  is  carefully  poured  a  quantity  of  water  until 
the  cylinder  is  nearly  full.  Upon  the  top  of  the 
water  is  then  floated  a  solution  of  potassium 
thiocyanate  in  alcohol,  and  the  whole  is  allowed  to  remain  undis- 
turbed. The  ferric  chloride  will  gradually  diffuse  up  into  the  water, 
and  the  dissolved  thiocyanate  will  diffuse  down,  and  at  the  point 
where  these  salts  meet  they  will  interact  chemically  upon  each, 
giving  rise  to  a  blood-red  coloured  solution,  which  will  appear  as 
a  ring  about  midway  down  the  cylinder. 

This  phenomenon  of  the  diffusion  of  dissolved  substances  is 
strictly  comparable  with  the  diffusion  of  gases,  although  in  the 
former  case  the  operation  proceeds  with  extreme  slowness.  The 
force  which  impels  the  molecules  of  dissolved  substances  to  diffuse 
is  the  osmotic  pressure  of  the  substance  in  solution. 


FIG.  26. 


160  Introductory  Outlines 

The  extension  of  the  gaseous  laws  into  the  domain  of  solutions 
necessitates  the  hypothesis  that  in  the  case  of  some  solutions  the 
molecules  of  the  dissolved  substance  unite  to  form  more  compli- 
cated molecular  associations  ;  while  in  other  cases  (including  those 
substances  which  are  electrolytes,  such  as  the  solutions  of  strong 
acids,  bases,  and  salts)  the  molecules  of  the  substances  undergo 
dissociation  into  their  ions.  For,  just  as  in  the  case  of  gases  where 
departures  from  the  strict  gaseous  laws  are  seen  to  take  place,  on 
account  of  the  dissociation  in  some  instances,  and  the  association 
in  others,  of  the  various  molecules,  so  it  is  believed  that  the  de- 
viations from  the  strict  continuity  of  the  ideal  gaseous  laws  into  the 
realm  of  solution  are  due  to  the  operation  of  similar  causes. 

CRYSTALLINE  FORMS. 

When  a  saturated  solution  of  a  solid  in  a  liquid  is  either  cooled 
or  allowed  to  evaporate,  the  dissolved  solid  begins  to  deposit  itself 
out  of  the  liquid,  and  it  does  so  in  most  cases  in  definite  geometric 
shapes,  termed  crystals.  (Solids  which  exhibit  no  crystalline 
structure  are  said  to  be  amorphous.} 

The  same  arrangement  of  molecules  into  geometric  forms  often 
takes  place  also  when  substances  in  a  state  of  fusion  (as  distin- 
guished from  solution)  pass  into  the  solid  condition,  as,  for  example, 
when  melted  sulphur,  or  mercury,  or  water  are  cooled  to  their  re- 
spective solidifying  points  ;  and  it  also  frequently  takes  place  when 
vapours  are  condensed  to  the  solid  state.  Speaking  generally,  the 
more  slowly  the  process  of  solidification  takes  place,  the  larger  and 
more  geometrically  perfect  will  be  the  crystals  that  are  formed. 

All  the  varieties  of  crystalline  forms,  both  naturally  occurring 
and  artificially  produced,  are  susceptible  of  classification  into  thirty- 
two  classes,*  based  upon  their  symmetrical  development  with  respect 
to  certain  imaginary  planes,  lines,  and  points,  called  respectively 
planes  of  symmetry ',  lines  of  symmetry,  and  centres  of  symmetry. 

Planes  of  symmetry  are  planes  cut  through  the  crystal  in  such  a 
direction  that  the  two  divided  portions  are  the  mirrored  reflections 
the  one  of  the  other,  the  mirror  being  the  plane  itself.  Crystals 
may  have  from  o  to  9  planes  of  symmetry ;  a  cube,  for  example,  has 
nine  such  planes. 

Axes  of  symmetry  are  imaginary  straight  lines  passing  through 
the  crystal  in  such  a  manner  that  when  the  crystal  is  rotated  upon 
one  of  them  there  will  be  a  complete  recurrence  of  similar  faces 

*  The  study  of  this  classification  belongs  to  the  science  of  crystallography, 
and  falls  outside  the  scope  of  a  general  chemical  text-book  ;  it  is  therefore  here 
treated  only  in  broadest  outline. 


Crystalline  Forms  161 

and  angles  at  least  once  before  an  entire  revolution  has  been  made. 
For  instance,  if  a  tube  is  rotated  upon  an  axes  passing  through  the 
centre  of  one  face  at  right  angles  to  the  face,  it  will  obviously 
present  the  same  appearance  four  times  during  a  complete  revolu- 
tion. In  thus  being  rotated  through  36o°*crystals  may  exhibit  this 
periodic  reappearance  of  the  same  aspect,  either  two,  three,  four, 
or  six  times,  and  the  axes  are  spoken  of  as  binary,  trigonal,  tetra- 
gonal, and  hexagonal  respectively.*  Crystals  may  possess  from 
o  to  13  axes  of  symmetry  ;  the  cube,  for  example,  has  thirteen  such 
axes,  viz.  six  binary,  four  trigonal,  and  three  tetragonal. 

Centres  of  symmetry.  A  crystal  has  a  centre  of  symmetry  when 
opposite  to  every  face  there  is  a  precisely  similar  face  parallel  to  it 
on  the  other  side  of  the  crystal. 

Based  upon  these  three  orders  of  symmetry  there  are  mathe- 
matically possible  thirty-two  classes  into  which  crystals  can  be 
ranged  ;  and  with  two  or  three  exceptions  only,  crystals  are  known 
belonging  to  each  class. 

Crystallographic  Systems.  These  thirty-two  classes  are  suscep- 
tible of  a  further  classification  into  the  following  six  systems,  based 
upon  the  relations  of  their  Crystallographic  axes  :  t— - 

I.  Cubic  (Regular  or  Isometric)  System.  Crystals  of  this  system  are  re- 
ferred to  three  axes  at  right  angles  to  each  other,  and  all  equal  in 
length.  The  system  includes  five  classes.  Of  these,  that  one  which 
exhibits  the  highest  degree  of  symmetry  (spoken  of  as  the  holohedral 
or  normal  class)  has  nine  planes  of  symmetry,  thirteen  axes  of  sym- 
metry, and  a  centre  of  symmetry. 

II.  Hexagonal  System.  Forms  of  this  system  are  referred  to  four  axes : 
three  are  equal  in  length  and  intersect  at  angles  of  120°,  while  the 
fourth,  known  as  the  principal  axis,  is  different  in  length  and  is 
vertical  to  the  others.  Twelve  classes  are  included  in  this  system.  J 
The  normal  or  holohedral  class  has  seven  planes  of  symmetry,  seven 
axes  of  symmetry,  and  a  centre  of  symmetry. 

III.  Tetragonal  System.     Crystals  belonging  to  this  system  are  referable  to 

three  axes  at  right  angles  to  each  other,  two  being  equal  in  length. 
The  system  includes  seven  classes,  the  hoiohedral  class  having  five 
planes,  five  axes,  and  a  centre  of  symmetry. 

IV.  Orthorhombic  {Rhombic  or  Prismatic]  System.     Crystals  are  referred  to 

three  axes  at  right  angles  to  each  other,  and  all  unequal  in  length. 
Three  classes  are  included  in  this  system,  the  holohedral  class  having 
three  planes,  three  axes,  and  a  centre  of  symmetry. 


*  Sometimes  the  terms  diad,  triad,  &c. ,  are  employed. 

f  Crystallographic  axes  do  not  necessarily  correspond  with  axes  of  symmetry, 
although  they  are  made  to  do  so  whenever  possible. 

J  Seven  of  these  classes  consist  of  rhombohedral  forms  of  this  system,  in 
which  the  principal  axis  is  a  trigonal  axis  of  symmetry  instead  of  being  one 
of  hexagonal  symmetry.  They  are  thus  regarded  as  hemihedral  (half  the 
number  of  faces)  modifications  of  the  hexagonal  form.  Some  crystallo- 
graphers  classify  them  as  a  separate  system  under  the  name  of  the  Rhombo- 
hedral system. 


162  Introductory  Outlines 

V.  Monosymmetric  (or  Monoclinic]  System.  In  this  system  the  forms  are 
referable  to  three  axes  of  unequal  lengths,  two  of  which  intersect  at 
an  acute  angle,  while  the  third  is  at  right  angles  to  the  other  two. 
Two  classes  belong  to  this  system,  the  holohedral  class  having  one 
plane,  one  axis,  and  a  centre  of  symmetry. 

VI.  Asymmetric  (or  Triclinic)  System.  Crystals  are  referred  to  three  axes 
of  unequal  lengths,  intersecting  one  another  at  oblique  angles.  Two 
classes  are  included  in  the  system,  the  holohedral  class  having  no 
planes  or  axes  of  symmetry,  but  a  centre  of  symmetry  only,  while  the 
second  class  has  no  element  of  symmetry  at  all. 

One  of  the  simplest  forms  in  each  of  these  six  systems  is  the  double 
pyramid  (see  Fig.  147  A).  Thus  there  is  the  tetragonal  pyramid, 
the  hexagonal  pyramid^  and  so  on.  In  the  case  of  the  isometric 
or  cubic  system  this  double  pyramid  is  called  the  Octahedron* 

Another  frequently  recurring  form  common  to  all  the  systems 
except  the  cubic,  is  that  of  the  prism,  giving  rise  to  tetragonal 
prisms^  hexagonal  prisms,  &c.f  Fig.  148  represents  a  group  of 
natural  crystals  in  the  form  of  hexagonal  prisms  terminating  in 
hexagonal  prisms.  Crystals,  whether  naturally  occurring  or  artifi- 
cially obtained,  very  seldom  exhibit  the  perfect  geometric  shape 
of  the  ideal  form,  but  usually  exhibit  more  or  less  distortion. 
Fig.  no  (the  left  crystal)  illustrates  distortion  in  an  orthorhombic 
pyramid.  Fig.  147  A  represents  an  octahedron  which  has  de- 
veloped into  the  perfect  ideal  form,  but  it  is  only  by  the  greatest 
care  in  regulating  their  growth  that  such  perfect  crystals  are 
obtained.  In  Fig.  147  B  is  seen  the  development  of  what  is  known 
as  twin  crystals.  It  very  often  happens  that  what  would  be  an 
edge  or  a  solid  angle  in  the  ideal  crystal  is  replaced  by  a  plane  or 
planes.  Such  variations  are  called  truncations.  Illustrations  of 
truncated  crystals  are  seen  in  Fig.  no.  In  the  right-hand  crystal 
both  apexes  are  truncated  by  planes  or  bases  ;  while  in  the  other 
crystal  one  apex  is  truncated  by  a  pyramid. 

Two  or  more  substances  which  crystallise  in  the  same  form  are 
said  to  be  isomorphous ;  and  on  the  other  hand  a  substance  which 
is  capable  of  crystallising  in  two  forms  which  do  not  belong  to  the 
same  system  is  termed  a  dimorphous  substance.  Thus  sulphur  is 
dimorphous  because  it  is  capable  of  crystallising  in  orthorhombic 
pyramids  and  in  monosymmetric  prisms.  Occasionally  a  dimor- 
phous substance  is  isomorphous  with  another  dimorphous  substance 
in  both  its  forms  ;  to  this  double  isomorphism  the  term  isodimor- 
phism  is  applied. 

*  The  double  pyramids  of  some  of  the  other  systems  are  also  octahedra,  in 
the  sense  that  they  possess  eight  faces,  but  in  modern  nomenclature  the  term 
Octahedron  is  reserved  exclusively  for  the  isometric  pyramid. 

t  It  will  be  obvious  that  a  description  of  a  crystal  merely  as  being  prismatic 
is  incomplete  without  reference  to  the  system  to  which  it  belongs. 


CHAPTER   XV 
THERMO-CHEMISTRY 

WE  have  seen  that  by  means  of  symbols  and  formulae  chemists 
express,  in  the  form  of  equations,  a  certain  amount  of  information 
respecting  chemical  changes:  thus  by  the  equation  C  +  O2=CO2 
there  are  conveyed  the  facts,  that  carbon  unites  with  oxygen  to 
form  carbon  dioxide,  that  12  grammes  of  carbon  combine  with  32 
grammes  of  oxygen,  yielding  44  grammes  of  carbon  dioxide,  and 
that  the  volume  of  the  gaseous  carbon  dioxide  obtained  is  the 
same  as  that  of  the  oxygen  taking  part  in  its  formation.  All  such 
equations  bear  upon  the  face  of  them  the  truth,  that  matter  can 
neither  be  destroyed  nor  created.  The  total  quantity  of  matter 
taking  part  in  the  action  is  unaltered  by  the  process,  although  it 
appears  in  altered  form  in  the  products  of  the  reaction. 

In  all  chemical  changes,  besides  matter,  energy  also  takes  a 
part ;  not  only  do  the  materials  concerned  undergo  rearrangement 
or  readjustment,  but  at  the  same  time  there  is  a  rearrangement  or 
readjustment  of  energy.  This  energy  change  is  not  expressed  by 
the  ordinary  symbolic  equation.  Thus  in  the  equation — 

SO3  +  H2O  =  H.,SO4 

the  fact  is  embodied  that  80  grammes  of  sulphur  trioxide  combine 
with  1 8  grammes  of  water  and  form  98  grammes  of  sulphuric  acid ; 
but  the  equation  takes  no  cognisance  of  the  fact,  that  when  these 
weights  of  these  two  substances  unite  to  form  98  grammes  of  sul- 
phuric acid  an  amount  of  energy,  in  the  form  of  heat,  is  disengaged 
that  would  raise  the  temperature  of  213  grammes  of  water  from  o° 
to  the  boiling-point. 

Similarly,  in  the  equation  2NC13=N2  +  3CI2  there  is  no  recogni- 
tion of  the  fact  that  during  this  change  an  enormous  amount  of 
energy  leaves  the  system  in  the  form  of  external  work  (over- 
coming the  atmospheric  pressure)  ;  in  other  words,  that  the  con- 
version of  nitrogen  trichloride  into  its  constituent  elements  is 
attended  with  the  most  violent  explosion. 

163 


164  Introductory  Outlines 

Energy,  like  matter,  can  neither  be  created  nor  destroyed,  but  as 
a  result  of  chemical  action  it  reappears  as  energy  in  another  form. 
Thus  it  may  appear  as  heat,  as  electrical  energy,  as  kinetic 
energy,  or  as  chemical  energy  ;  and  just  as  the  total  amount  of 
matter  taking  part  in  a  chemical  change  reappears  in  altered  form 
in  the  products  of  the  change,  so  the  disappearance  of  energy  in 
any  of  its  forms  gives  rise  to  the  reappearance  of  a  proportionate 
amount  of  energy  in  another  form.  This  is  the  law  of  the  conserva- 
tion of  energy )  which  may  be  thus  stated  :  *  "  The  total  energy  of  any 
material  system  is  a  quantity  which  can  neither  be  increased  nor 
diminished  by  any  action  between  the  parts  of  the  system,  although 
it  may  be  transformed  into  any  of  the  forms  of  which  energy  is 
susceptible? 

Chemical  energy,  or  that  form  of  energy  that  is  set  free  during 
chemical  processes,  cannot  be  measured  by  any  direct  method. 
This  energy,  however,  is  generally  transformed,  during  chemi- 
cal change,  into  heat,  and  may  therefore  be  measured  by,  and 
expressed  in,  heat  units.  Thermo-chemistry  may  therefore  be 
defined  as  the  science  of  the  thermal  changes  which  accompany 
chemical  changes. 

All  matter  is  regarded  as  containing  a  certain  amount  of  energy 
in  some  form,  and  the  purpose  of  thermo-chemistry  is,  by  measur- 
ing the  thermal  disturbance  that  is  conditioned  by  a  chemical 
change,  to  ascertain  the  difference  between  the  amount  of  energy 
contained  in  a  system  before  and  after  such  a  change. 

If  all  the  energy  of  a  system  in  its  original  state  (i.e.  before  the 
chemical  change  takes  place)  that  undergoes  transformation  into 
other  forms  of  energy  passes  into  heat  ;  if  none  of  it  leaves  the 
system  as  energy  in  some  other  form,  and  thereby  escapes  mea- 
surement ;  then  the  difference  between  the  amount  of  energy 
contained  in  the  system  in  its  original  and  its  final  state  may  be 
ascertained.  It  by  no  means  follows,  however,  that  this  represents 
the  chemical  energy  alone  ;  it  has  already  been  explained  that 
chemical  changes  are  always  attended  by  physical  changes,  such 
as  change  of  volume,  of  physical  state,  and  so  on,  and  we  have 
also  learned  that  such  physical  changes  are  likewise  accompanied  by 
thermal  changes;  the  problem,  therefore,  is  often  a  complicated  one, 
and  it  is  not  always  possible  to  differentiate  between  the  chemical 
and  the  physical  causes  that  may  be  operating  simultaneously,  and 
to  decide  what  share  of  the  final  result  is  due  to  the  chemical  phase 
*  Clerk  Maxwell,  "  Matter  and  Motion." 


Thermo-  Chem  is  try  165 

of  the  change,  and  what  to  the  physical  change  that  simultaneously 
takes  place. 

As  an  illustration  of  the  complex  nature  of  chemical  reactions 
when  considered  from  a  thermal  standpoint,  and  of  the  disturbing 
effect  of  the  accompanying  physical  changes,  we  may  take  the  case 
of  the  action  of  aqueous  hydrochloric  acid  upon  crystallised  sodium 
sulphate,  Na2SO4,10H2O— 


Na2SO4,10H2 

The  chemical  action  here  consists  of  (i)  the  decomposition  of 
sodium  sulphate,  (2)  the  decomposition  of  hydrochloric  acid,  (3) 
the  formation  of  sodium  chloride,  (4)  the  formation  of  sulphuric 
acid.  Heat  is  absorbed  by  the  first  two  portions  of  the  action,  and 
heat  is  evolved  by  the  other  two.  The  physical  changes  include 
the  passage  of  ten  molecules  of  water  of  crystallisation  (i.e.  solid 
water)  into  liquid  water,  and  the  solution  of  sodium  chloride  in 
water.  These  changes  are  attended  with  absorption  of  heat,  and 
the  net  result  of  the  entire  change  is  the  disappearance  of  a  con- 
siderable amount  of  heat,  that  is  to  say,  the  thermal  value  of  the 
reaction  is  a  negative  quantity. 

The  methods  adopted  in  order  to  express  thermo-chemical 
reactions  are  quite  simple.  The  ordinary  chemical  symbols  and 
formulae  are  used,  and  represent,  in  all  cases,  quantities  in  grammes 
corresponding  to  the  formula-weights  of  the  substances.  Thus  Cl 
represents  35.5  grammes  of  chlorine  ;  H2O  stands  for  18  grammes 
of  water,  and  so  on.  The  chemical  equation  is  followed  by  a 
number  representing  the  quantity  of  heat,  expressed  in  heat  units, 
which  is  either  produced  or  which  disappears  as  a  result  of  the 
change.  The  unit  of  heat  is  the  calorie,  or  the  quantity  of  heat 
that  is  capable  of  raising  the  temperature  of  I  gramme  of  water  from 
o°  to  i°.  Sometimes  the  unit  employed  is  the  quantity  of  heat 
required  to  raise  I  gramme  of  water  from  o°  to  100°,  and  this  unit 
(which  is  100  times  greater  than  the  calorie)  is  indicated  usually 
by  the  letter  K.  When  heat  is  produced  by  a  chemical  change, 
the  sign  +  is  placed  in  front  of  the  number  of  units,  and  when 
heat  disappears,  the  fact  is  indicated  by  the  sign  -  . 

Thus  the  equation  — 


=  2HC1  +44,000  cal. 
or  H2+C12=2HC1  +  440  K, 

means   that   when   2   grammes    of   hydrogen   combine    with   71 


1 66  Introductory  Outlines 

grammes  of  chlorine  to  form  gaseous  hydrochloric  acid,  heat  is 
disengaged  to  the  amount  of  44,000  calories,  or  440  of  the  larger 
units,  K.  Or,  in  other  words,  that  when  these  quantities  of  these 
substances  combine,  an  amount  of  energy  is  lost  to  the  system, 
represented  by  44,000  calories.  Therefore  the  energy  possessed 
by  2  grammes  of  hydrogen  and  71  grammes  of  chlorine  is  greater 
than  that  possessed  by  73  grammes  of  hydrochloric  acid  gas  by  an 
amount  which  is  represented  by  44,000  gramme-units  of  heat. 
Hence  the  equation  may  be  written — 

2H  Cl  =  H2  +  C12  -  44,000  cal. 

which  signifies  that  when  73  grammes  of  gaseous  hydrochloric  acid 
are  decomposed  into  chlorine  and  hydrogen,  it  is  necessary  to 
supply  an  amount  of  energy  equal  to  44,000  calories. 

In  order  to  indicate  the  state  of  aggregation  of  the  different  sub- 
stances, the  method  introduced  by  Ostwald  consists  in  the  use  of 
different  type,  thick  type  being  employed  to  denote  solids,  ordinary 
type  indicating  liquids,  and  italics  signifying  gases,  thus — 

C  +  <92 = C<92 + 97,000  cal. 

means  that  the  total  energies  of  12  grammes  of  solid  carbon  and 
32  grammes  of  gaseous  oxygen  is  greater  than  the  energy  pos- 
sessed by  44  grammes  of  gaseous  carbon  dioxide  by  an  amount 
equivalent  to  97,000  calories. 
Or,  again,  the  equation — 

S03  +  H2O  =  H2SO4+2i,32o  cal. 

signifies  that  80  grammes  of  solid  sulphur  trioxide  unites  with  18 
grammes  of  liquid  water  and  forms  98  grammes  of  liquid  sulphuric 
acid,  with  the  liberation  of  21,300  gramme-units  of  heat. 

Similarly,  the  heat  evolved  by  the  passage  of  water  into  ice,  and 
the  heat  that  disappears  when  water  passes  into  steam,  may  be 
expressed  by  the  equations — 

H2O  =  H20  + 1440  cal. 
H2O  =  ^0-9670  cal. 

when  water  takes  a  direct  part  in  the  chemical  change,  as,  for 
example,  in  the  action  of  sulphur  trioxide  and  water  already  quoted, 
the  formula  represents  a  gramme-molecule  just  as  in  all  other 


Thermo-  Chemistry  1  67 

cases  ;  but  where  the  presence  of  a  large  quantity  of  water  affects  the 
thermal  result  of  the  chemical  change,  by  exerting,  for  example,  a 
solvent  action,  the  symbol  Aq  is  employed  to  signify  that  the  pre- 
sence of  the  water  is  considered  in  the  thermal  expression. 
Thus  the  expression  — 

/^r+Aq=HBrAq+  19,900  cal. 

signifies  that  when  81  grammes  of  gaseous  hydrobromic  acid  are 
dissolved  in  a  large  excess  of  water,  19,900  calories  are  evolved. 
Again,  the  equation  — 

//2+^2  +  Aq  =  2HBrAq  +  64,ooo  cal. 

means  that  when  160  grammes  of  gaseous  bromine  combine  with 
2  grammes  of  hydrogen,  and  the  product  is  dissolved  in  an  excess 
of  water  (i.e.  such  a  quantity  of  water  that  no  thermal  change  is 
produced  by  the  addition  of  any  further  quantity),  64,000  calories 
are  disengaged.  Of  this  64,000  calories,  19,900x2  =  39,800  are 
due  to  the  solution  of  the  twice  81  grammes  of  hydrobromic  acid, 
and  the  difference,  viz.,  24,000  calories,  represent  the  heat  produced 
by  the  combination  of  2  grammes  of  hydrogen  with  160  grammes 
of  bromine. 

If  water  is  formed  as  one  of  the  products  of  the  chemical  reaction 
taking  place  in  the  case  of  substances  in  aqueous  solution,  such  as 
when  a  solution  of  hydrochloric  acid  is  added  to  a  solution  ot 
sodium  hydroxide,  HCl  +  NaHO  =  NaCl  +  H2O,  as  the  water  so 
produced  simply  mixes  with  the  water  in  which  the  materials  are 
dissolved,  without  producing  any  thermal  effects  by  so  doing,  it  is 
usually  neglected  in  energy  equations  ;  although,  as  already  stated 
(page  109),  when  explained  from  the  standpoint  of  the  ionic  theory, 
the  heat  of  neutralisation  is  here  due  to  the  formation  of  mole- 
cules of  water  by  the  union  of  H'ions  with  HO'  ions.  Thus  the 
above  action  may  be  expressed  — 

HClAq+NaHOAq  =  NaClAq+  13,736  cal. 

The  heat  that  is  produced,  or  that  disappears,  in  a  chemical 
change  which  results  in  the  formation  of  a  particular  compound 
is  termed  the  heat  of  formation  of  that  compound.  Thus  in  the 
equation  — 

7/2  +C/2  =  2^7C/+  44,000  cal. 


the  heat  of  formation  of  73  grammes  of  hydrochloric  acid  is  44,000 


1  68  Introductory  Outlines 

thermal  units.  This  number,  however,  is  in  reality  the  algebraic  sum 
of  three  quantities.  It  does  not  express  merely  the  heat  developed 
by  the  simple  union  of  chlorine  and  hydrogen.  The  chemical 
change  expressed  by  the  equation  consists  in  reality  of  three 
operations  — 

(i.)  H2  =  H  +  H.    (2.)C12=C1  +  C1.    (3.)C1  +  C1  +  H  +  H  =  2HC1. 

Each  of  these  operations  represents  a  distinct  thermal  effect  ;  in 
Nos.  (i)  and  (2)  heat  is  absorbed,  in  No.  (3)  heat  is  evolved,  and 
calling  these  values  h-^  h^  and  /&3,  we  have  as  the  net  result 
^3  ~  (h\  +  ^2)  =  44,000  cal. 

The  number  of  heat-units,  therefore,  which  expresses  the  heat  of 
formation  of  hydrochloric  acid  is  the  heat  produced  by  the  union  of 
two  atoms  of  hydrogen  with  two  atoms  of  chlorine,  minus  the  heat 
absorbed  in  the  decomposition  of  one  hydrogen  and  one  chlorine 
molecule. 

Compounds  such  as  hydrochloric  acid,  in  the  formation  of  which 
heat  is  developed,  are  termed  exothermic  compounds,  the  reaction 
by  which  they  are  produced  being  an  exothermic  change  ;  com- 
pounds, on  the  other  hand,  whose  heats  of  formation  are  expressed 
by  a  negative  sign,  that  is,  in  whose  formation  heat  disappears,  are 
distinguished  as  endothermic  compounds,  and  the  reactions  by 
which  they  are  formed  are  endothermic  reactions. 


Thus  C  +  S2  =  CS2-  19,600  cal., 

signifies  that  in  the  formation  of  carbon  disuiphide  heat  is  absorbed, 
and  the  compound  is  therefore  an  endothermic  compound. 

Thermo-chemical  determinations  are  made  by  means  of  instru- 
ments termed  calorimeters.  These  are  of  great  variety,  although 
the  principle  involved  is  the  same.  The  chemical  reaction  is  caused 
to  take  place  under  such  circumstances,  that  the  whole  of  the  heat 
that  is  liberated  shall  be  communicated  to  a  known  volume  of 
water,  at  a  known  temperature.* 

Direct  determinations  of  the  thermal  value  of  chemical  changes 
have  hitherto  been  made  in  only  a  limited  number  of  comparatively 
simple  cases  ;  it  is  possible,  however,  from  a  few  known  data,  to  cal- 
culate the  thermal  values  of  a  number  of  changes  which  cannot  be 
directly  measured.  This  depends  upon  the  fundamental  principle 

*  For  descriptions  of  the  various  calorimeters,  see  treatises  on  Physics. 


Therm  o-  Chem  is  try  1 69 

of  thermo-chemistry,  which  is  itself  the  corollary  of  the  law  of  the 
conservation  of  energy,  and  which  was  first  experimentally  proved 
by  Hess  (1840).  This  principle,  which  is  sometimes  termed  the  law 
of  constant  heat  consummation,  or  the  law  of  equivalence  of  heat 
and  chemical  change^  may  be  thus  stated  :  The  amount  of  heat 
that  is  liberated  or  absorbed,  during  a  chemical  process,  is  de- 
pendent solely  upon  the  initial  and  final  states  of  the  system,  and  is 
independent  of  the  intermediate  stages.  The  following  examples 
will  serve  to  explain  the  application  of  the  principle  : — 

1.  Let  us  suppose  it  is  desired  to  find  the  heat  of  formation  of 
carbon  monoxide,  the  data  at  our  disposal  being  (i)  the  heat  pro- 
duced when  carbon  unites  with  oxygen  to  form  carbon  dioxide  ;  and 
(2)  the  heat  formed  by  the   combustion  of  carbon  monoxide  to 
carbon  dioxide.     The  thermal  equations  are — 

(1)  C +  02  =  C02+ 97,000  cal. 

(2)  2CO+  O2=2CO2+  136,000  cal. 

Halving  the  second  equation,  in  order  to  get  the  heat  produced 
in  the  formation  of  44  grammes  of  carbon  dioxide  (i.e.  the  same 
weight  as  in  the  first),  we  may  represent  the  equation  as — 

CO  +  O  =  COZ  +  68,000  cal.* 

The  difference  between  the  two  values  97,000  and  68,000  will  be 
the  heat  of  formation  of  carbon  monoxide,  therefore  we  get  the 
equation — 

G  +  O  =  CO  +  29,000  cal. 

2.  The  compound,  methane  (marsh  gas),  CH4,  cannot  be  formed 
by  the  direct  union  of  its  elements,  but  its  heat  of  formation  can 
be  calculated  by  the  application  of  this  principle.      The  data  in 
this  case  are  the  ascertained  heats  of  formation  of  carbon  dioxide 

*  It  must  be  remembered  that  this  equation  does  not  express  the  whole 
truth  :  as  it  here  stands  it  would  imply  that  68,000  calories  represent  the  heat 
formed  by  the  simple  chemical  union  of  28  grammes  of  carbon  monoxide  with 
16  grammes  of  oxygen.  In  reality  this  number  is  half  the  sum  of  the  two 
values,  namely,  the  heat  of  combination  of  56  grammes  of  carbon  monoxide 
with  32  grammes  of  oxygen,  minus  the  heat  absorbed  by  the  decomposition  of 
32  grammes  of  oxygen  molecules  into  their  constituent  atoms.  The  oxygen 
atom  does  not  exist  alone,  and  whenever  free  oxygen  takes  part  in  a  chemical 
change  the  molecules  of  the  element  are  first  separated  into  their  atoms. 


170  Introductory  Outlines 

and  of  water,  and  the  heat  produced  by  the  combustion  of  marsh 
gas,  the  thermal  equations  being — 

(1)  C  +  <92=  C6>2  +  97,000  cal. 

(2)  2/y?+02  =  2/720+  136,800  cal. 

(3)  C^4  +  2<92=C<92  +  2#26>  +  2i2,ooocal. 

The  difference  between  the  thermal  value  of  the  last  process  and 
the  sum  of  the  first  and  second  represents  the  heat  of  formation  of 
marsh  gas — 

97,000+136,800  —  212,000  =  21,800, 

hence  we  get  the  expression — 

C  +  2#"2  =  C//"4  +  2 1 ,800  cal. 


PART   II 

THE  STUDY  OF  FOUR  TYPICAL  ELEMENTS 
HYDROGEN-OXYGEN-NITROGEN-CARBON 

AND  THEIR   MORE  IMPORTANT  COMPOUNDS 

CHAPTER    I 

HYDROGEN 

Symbol,  H.     Atomic  weight  =  1.008.     Molecular  weight =2.016. 
Density  =  i.  008. 

History. — The  existence  of  hydrogen  as  an  individual  sub- 
stance was  first  established  by  Cavendish  (1766),  who  applied  to  it 
the  name  inflammable  air.  He  obtained  the  gas  by  acting  upon 
certain  metals,  as  iron,  tin,  and  zinc,  with  either  sulphuric  or  hydro- 
chloric acid. 

Occurrence. — In  the  free  state  hydrogen  occurs  only  in  small 
quantities  upon  the  earth.  It  is  evolved  with  other  volcanic  gases, 
and  is  present  in  the  gases  which  escape  from  petroleum  wells. 
It  is  evolved  also  during  the  fermentation  and  decomposition  of 
certain  organic  compounds,  and  is  therefore  present  in  the  breath 
and  the  intestinal  gases  of  animals.  From  these  sources  it  finds 
its  way  into  the  atmosphere,  where  it  is  present  to  the  extent  of 
about  .04  volumes  in  1000  volumes  of  air.  Hydrogen  Has  also 
been  found  in  many  specimens  of  meteoric  iron,  and  also  in 
certain  rocks,  where  it  is  present  as  occluded  gas. 

Hydrogen  in  the  uncombined  state  exists  in  enormous  masses 
upon  the  sun,  and  is  present  in  certain  stars  and  nebulae.  The 
so-called  prominences  which  are  seen  projecting  from  the  sun's 
disk  to  a  distance  of  many  thousands  of  miles,  and  which  were 


172  Inorganic  Chemistry 

first  observed  during  solar  eclipses,  consist  of  vast  masses  of  in- 
candescent hydrogen. 

In  combination  with  other  elements  hydrogen  is  extremely 
abundant ;  its  commonest  compound  is  water,  which  consists  of 
one  part  by  weight  of  this  element  combined  with  eight  parts  of 
oxygen.  In  combination  with  chlorine,  as  hydrochloric  acid,  with 
carbon  as  marsh  gas,  and  with  sulphur  as  sulphuretted  hydrogen, 
this  element  also  occurs  in  large  quantities.  All  known  acids 


FIG.  27. 

contain  hydrogen  as  one  of  their  constituents,  and  it  is  present  in 
almost  all  organic  compounds. 

Modes  of  Formation. — (i.)  Hydrogen  may  be  obtained  from 
water  by  the  action  of  various  metals  upon  that  compound  under 
certain  conditions.  The  metals  sodium  and  potassium  will  decom- 
pose water  at  the  ordinary  temperatures  ;  when,  therefore,  a  frag- 
ment of  either  of  these  metals  is  thrown  upon  water,  the  latter  is 
decomposed  and  hydrogen  set  free  : — 

=  NaHO  +  H. 


Hydrogen 

The  metals,  being  lighter  than  water,  float  upon  its  surface,  and, 
owing  to  the  heat  of  the  reaction,  melt  and  roll  about  upon  the  liquid 
as  molten  globules.  With  potassium,  the  temperature  developed 
is  sufficiently  high  to  cause  the  hydrogen  to  inflame,  and  it  burns 
with  a  flame  coloured  violet  by  the  vapour  of  the  metal.  The 
hydroxide  of  the  metal,  which  is  the  second  product  of  the  action, 
dissolves  in  the  excess  of  water,  rendering  the  liquid  alkaline. 
The  alkalinity  of  the  solution  may  be  made  evident  by  the  addition  of 
a  reddened  solution  of  litmus,  which  will  be  turned  blue  by  the  alkali. 

In  order  to  collect  the  hydrogen  evolved  by  the  action  of  sodium 
upon  water,  the  metal  is  placed  in  a  short  piece  of  lead  tube  closed 
at  one  end,  which  causes  it  to  sink  in  the  liquid,  and  an  inverted 
glass  cylinder  filled  with  water  is  placed  over  it,  as  shown  in  Fig.  27. 
The  evolved  hydrogen  then  rises  as  a  stream  of  bubbles  into  the 
cylinder  and  displaces  the  water.* 


FIG.  28. 

(2.)  Water  may  be  readily  decomposed  at  the  boiling-point, 
by  means  of  zinc,  if  the  metal  be  previously  coated  with  a  thin 
film  of  copper  by  immersion  in  a  dilute  solution  of  copper  sul- 
phate. When  this  copper-coated  zinc  (known  as  zinc  -copper 
couple)  is  heated  in  a  small  flask  filled  with  water,  and  provided 
with  a  delivery  tube,  the  oxygen  of  the  water  combines  with  the 
zinc  forming  zinc  oxide,  and  hydrogen  is  evolved,  which  may  be 
collected  over  water  at  the  pneumatic  trough  :  *— 


*  For  detailed  description  of  these  experiments,  see  Newth's  "Chemical 
Lecture  Experiments,"  p.  2. 


174 


Inorganic  Chemistry 


(3.)  At  a  still  higher  temperature,  water  in  the  state  of  steam  can 
be  readily  decomposed  by  the  metal  magnesium,  magnesium  oxide 
being  formed  and  hydrogen  liberated  :  — 


For  this  purpose  the  magnesium  is  strongly  heated  in  a  glass 
bulb  (Fig.  28),  while  steam  from  a  small  boiler  is  passed  over  it. 
As  the  temperature  of  the  metal  approaches  a  red  heat  it  bursts 
into  flame,  and  the  issuing  hydrogen  may  be  ignited  as  it  escapes 
from  the  end  of  the  tube. 

(4.)  If  iron  be  heated  to  bright  redness  and  steam  be  passed 
over  it,  the  water  is  decomposed,  the  oxygen  uniting  with  the  iron 


FIG.  29. 

to  form  an  oxide  known  as  triferric  tetroxide,  or  magnetic  oxide  of 
iron,  thus — 

3Fe  +  4H2O  =  Fe3O4  +  4H2. 

This  method  is  employed  on  a  large  scale  for  the  preparation  of 
hydrogen  for  commercial  purposes.  Iron  borings  or  turnings  are 
packed  into  an  iron  tube,  which  is  strongly  heated  in  a  furnace, 
and  steam  from  a  boiler  is  passed  through  the  tube. 

(5.)  For  laboratory  purposes  hydrogen  is  most  conveniently  pre- 
pared by  the  action  of  dilute  sulphuric  acid  upon  zinc  : — 

Zn  +  H2SO4=ZnSO4+H2. 
For  this  purpose  granulated  zinc  (i.e.  zinc  which  has  been  melted 


Hydrogen  175 

and  poured  into  water)  is  placed  in  a  two-necked  Woulf's  bottle 
(Fig.  29),  and  a  quantity  of  sulphuric  acid,  previously  diluted  with 
six  times  its  volume  of  water,  is  introduced  by  means  of  the  funnel. 
A  brisk  action  sets  in,  and  hydrogen  is  rapidly  disengaged.  After 
the  lapse  of  a  few  minutes  the  air  within  the  apparatus  will  be 
swept  out  by  the  hydrogen,  when  the  gas  may  be  collected  over 
water  in  the  pneumatic  trough. 

The  hydrogen  so  obtained  is  never  absolutely  pure  ;  it  is  liable  to  contain 
traces  of  arsenic  hydride,  hydrogen  sulphide,  hydrogen  phosphide,  oxides  of 
nitrogen,  and  nitrogen.  The  nitrogen  is  derived  from  the  air,  which  finds 
its  way  through  joints  in  the  apparatus,  and  also  from  air  dissolved  in  the 
acid.  There  is  no  known  process  for  removing  this  impurity.  The  other 
gases  are  due  to  impurities  in  the  zinc  and  the  sulphuric  acid,  and  can  be 
removed,  if  required,  by  passing  the  hydrogen  through  a  series  of  tubes  con- 
taining absorbents  (see  page  210). 

Absolutely  pure  sulphuric  acid,  even  when  diluted  with  water, 
"las  no  action  upon  perfectly  pure  zinc. 

Scrap  iron  may  be  substituted  for  zinc,  but  the  hydrogen  so 

obtained  is  much  less  pure,  and  is  accompanied  by  compounds  of 

carbon  (derived  from  the  carbon  in  the  iron),  which  impart  to  the 

gas  an  unpleasant  smell  ;  the  reaction  in  this  case  is  the  following  :  — 

Fe  +  H2SO4  =  FeSO4+H2. 

Hydrochloric  acid  can  be  employed  in  place  of  sulphuric  acid 
with  either  zinc  or  iron,  the  reaction  then  being:  — 
Zn  +  2HCl  =  ZnCl2+H2. 

These  actions  of  acids  upon  metals  when  expressed  in  the  form  of 
ionic  equations  will  each  be  seen  to  consist  simply  of  the  transference 
of  two  positive  atomic  charges  from  two  hydrogen  ions  to  an  atom 
of  the  metal,whereby  the  latter  is  converted  into  a  divalent  ion,  thus  — 


(6.)  Hydrogen  in  a  high  degree  of  purity  is  conveniently  prepared 
in  small  quantity  by  the  electrolysis  of  water  acidulated  with  sul- 
phuric acid  (see  page  207). 

(7.)  Hydrogen  is  disengaged  when  certain  metals,  such  as  zinc, 
iron,  and  aluminium,  are  boiled  with  an  aqueous  solution  of  potas- 
sium or  sodium  hydroxide.  Thus,  in  the  case  of  zinc,  when  this 
metal  in  the  form  of  filings  is  boiled  with  a  solution  of  potassium 
hydroxide,  hydrogen  is  evolved,  and  a  compound  of  zinc,  potassium, 
and  oxygen  remains  in  solution,  namely,  potassium  zinc  oxide  (or 
potassium  zincate),  thus  :  — 


(8.)  Hydrogen  is  also  obtained  by  heating  alkaline  oxalates,  or 


176  Inorganic  Chemistry 

formates,  with  either  potassium  or  sodium  hydroxide,  with  the 
simultaneous  formation  of  an  alkaline  carbonate  ;  thus  with  sodium 
oxalate : — 


Properties. — Hydrogen  is  a  colourless  gas,  and  has  neither 
taste  nor  smell.  It  is  the  lightest  known  substance,  being  14.3875 
times  lighter  than  air.  Its  specific  gravity  is  0.0695  (air=i)- 
One  litre  of  the  gas  at  o°  C.,  and  under  a  pressure  of  760  mm.  of 
mercury  (i.e.  the  standard  temperature  and  pressure)  weighs 
0.089873  gramme  ;  or  I  gramme  of  hydrogen  at  the  standard  tem- 
perature and  pressure  occupies  11.127  litres. 

On  account  of  its  extreme  lightness,  hydrogen  may  be  poured 


FIG.  30. 

upwards  from  one  vessel  to  another.  If  a  large  beaker  be  sus- 
pended mouth  downwards  from  the  arm  of  a  balance  and  counter- 
poised, and  the  contents  of  a  jar  of  hydrogen  be  poured  upwards 
into  the  beaker,  the  equilibrium  of  the  system  will  be  disturbed, 
and  the  arm  carrying  the  beaker  will  rise. 

The  lightness  of  hydrogen  can  also  be  shown  by  causing  a 
stream  of  the  gas  to  issue  from  a  tube  placed  in  such  a  position 
that  its  shadow  is  cast  upon  a  white  screen  by  means  of  a  powerful 
electric  light.  When  the  gas  is  streaming  from  the  tube,  its  up- 
ward rush  will  be  visible  upon  the  screen  as  a  distinct  shadow, 
caused  by  the  difference  between  the  refractive  power  of  air  and 
hydrogen  (Fig.  30). 


Hydrogen 


177 


Hydrogen  is  inflammable  and  burns  with  a  non-luminous  flame, 
the  temperature  of  which  is  very  high.  The  product  of  the  com- 
bustion of  hydrogen  is  water,  and  if  a  jet  of  the  gas  be  burned 
beneath  the  apparatus  seen  in  Fig.  31,  considerable  quantities  of 
water  may  be  collected  in  the  bulb.  In  the  act  of  combustion  the 
hydrogen  combines  with  the  oxygen  of  the  air,  forming  the  oxide 
of  hydrogen,  namely,  water:*  — 


If  hydrogen  be  mixed  with  the  requisite  quantity  of  air,  or  oxygen, 
and  a  light  applied  to  the  mixture,  the 
combination  of  the  two  gases  takes 
place  instantly  with  a  violent  explo- 
sion ;  t  hence  the  necessity  of  care- 
fully expelling  all  the  air  from  the 
apparatus  in  which  hydrogen  is  being 
generated  before  applying  a  flame  to 
the  issuing  gas. 

Hydrogen  will  not  support  the  com- 
bustion of  ordinary  combustibles  ;  thus, 
if  a  burning  taper  be  thrust  into  a  jar  of 
the  gas,  the  hydrogen  itself  will  be 
ignited  at  the  mouth  of  the  jar,  which 
must  be  held  in  an  inverted  position, 
but  the  taper  will  be  extinguished  ;  on 
withdrawing  the  taper  it  may  be  re-  | 
ignited  by  the  burning  hydrogen. 

Although  hydrogen  is  not  poisonous, 
it  is  incapable  of  supporting  animal  life 
owing  simply  to  the  exclusion  of  oxygen. 
When  mixed  with  air  and  inhaled,  it  raises  the  pitch  of  the  voice 


FIG.  31. 


*  From  this  fact  the  name  Hydrogen  (signifying  the  water  producer]  is  derived. 

f  Baker  has  recently  shown  (four.  Chem.  Soc. ,  April  1902)  that  if  the  tv/o 
gases  are  perfectly  pure  and  dry,  they  may  be  strongly  heated  without  uniting. 
In  these  experiments  a  coil  of  silver  wire  suspended  in  the  gases  was  heated  by 
means  of  an  electric  current  until  the  silver  melted,  that  is,  above  1000°;  but 
no  chemical  union  of  the  oxygen  and  hydrogen  took  place,  although  the  ordi- 
nary temperature  of  explosion  is  615°  (V.  Meyer).  Baker  has  also  shown  that  if 
a  mixture  of  these  two  gases,  which  has  not  been  specially  dried,  be  exposed  to 
sunlight,  combination  slowly  takes  place;  whereas  with  the  perfectly  dry 
gases  no  measurable  combination  occurs. 

M 


178  Inorganic  Chemistry 

almost  to  a  falsetto.  The  same  effect  may  be  seen  by  sounding  a 
pitch-pipe,  or  organ-pipe,  by  means  of  a  stream  of  hydrogen 
instead  of  ordinary  air,  when  it  will  be  noticed  that  the  note  given 
out  is  greatly  raised  in  pitch. 

Hydrogen  is  very  slightly  soluble  in  water.  It  was  formerly 
believed  that  this  gas  formed  an  exception  to  the  rule  that  the 
solubility  of  gases  in  water  diminishes  with  rise  of  temperature, 
and  it  was  supposed  that  the  solubility  of  hydrogen  was  constant 
between  the  temperatures  o°  and  25°.  More  recent  experiments 
have  shown  that  this  is  not  the  case.  The  solubility  of  this  gas,  as 
determined  by  W.  Timofejeff(i89o),  is  seen  in  the  table  on  p.  143. 

Hydrogen  was  first  liquefied  on  May  10,  1898,  by  Dewar.  Prior 
to  this  time  it  had  never  been  obtained  as  a  coherent  or  static 
liquid — that  is,  a  liquid  with  a  meniscus — although  momentary 
indications  of  its  liquefaction  had  been  obtained  by  Olszewski  as 
far  back  as  1895.  The  critical  temperature  of  hydrogen  (namely, 

—  238°)  being  below   the   lowest  point  obtainable  by  the   rapid 
ebullition  of  liquid  oxygen  or  air,  no  external  refrigerating  agent 
is  available  which  is  capable  of  cooling  the  gas  below  its  critical 
point,  and  therefore  of  causing  its  liquefaction.     By  an  extension 
of  the  principle  of  self-cooling  explained  on  p.  76,  however,  namely, 
by  causing  a  jet  of  the  gas  previously  cooled  to  —205°  to  continu- 
ously escape  from  a  fine  orifice  under  a  pressure  of  180  atmos- 
pheres, Dewar  has  succeeded  in  collecting  considerable  quantities 
of   liquid    hydrogen    in    specially   constructed  vacuum -jacketed 
vessels. 

Liquid  hydrogen  is  clear  and  colourless  as  water,  thus  disposing 
of  the  theory  once  advocated  that  if  obtained  in  the  liquid  state 
hydrogen  would  be  found  to  exhibit  metallic  properties.  The 
boiling-point  of  the  liquid  is  —253°  (Dewar),  at  which  temperature 
air  is  immediately  solidified.  Thus,  if  a  tube  sealed  at  one  end, 
but  freely  open  to  the  air  at  the  other,  be  immersed  in  liquid 
hydrogen,  the  cooled  end  of  the  tube  quickly  becomes  filled  with 
solidified  air.  Similarly,  oxygen  is  frozen  to  a  pale-blue  solid. 

The  specific  gravity  of  liquid  hydrogen  is  about  0.07  ;  that  is  to 
say,  it  is  only  about  ^th  the  density  of  water,  or  about  14  c.c.  of  the 
liquid  weighs  only  I  gramme.  By  its  own  rapid  evaporation  liquid 
hydrogen  has  been  frozen  to  a  white  solid  mass,  which  melts  at 

—  257°  ;  and  by  the  rapid  evaporation  of  this  solid  a  temperature 
of  —260°  has  been  obtained,  which  is  the  lowest  degree  of  cold 
ever  reached.     By  means  of  liquid  hydrogen  as  a  refrigerating 


Hydrogen  1 79 

agent,  all  the  known  gases,  have  been  condensed  to  the  liquid 
state. 

Occluded  Hydrogen. — Certain  metals,  such  as  iron,  platinum, 
and  notably  palladium,  possess  the  property  when  heated  of 
absorbing  a  large  quantity  of  hydrogen,  and  of  retaining  it  when 
cold.  Graham  found  that  at  a  red  heat  palladium  absorbed,  or 
occluded,  about  900  times  its  own  volume  of  hydrogen,  while 
even  at  ordinary  temperatures  it  was  able  to  absorb  as  much 
as  376  times  its  volume.*  Graham  believed  that  the  hydrogen 
so  occluded  assumed  the  solid  form,  and  was  alloyed  with  the 
palladium,  and  to  denote  the  metallic  nature  of  the  gas  he  gave 
to  it  the  name  hydrogenium.  From  later  experiments  of  Troost 
and  Hautefeuille,  it  seems  probable  that  a  definite  compound  of 
hydrogen  and  palladium  exists,  of  the  composition  of  Pd2H. 

After  its  absorption  of  hydrogen  the  metal  presents  the  same 
appearance  as  before,  although  some  of  its  physical  properties 
have  become  slightly  modified  ;  thus  it  is  more  magnetic  than 
ordinary  palladium,  and  its  electric  conductivity  is  considerably 
reduced. 

In  view  of  our  present  knowledge  of  the  entire  absence  of  any 
metallic  characters  in  liquid  or  solid  hydrogen  (gained,  however, 
entirely  since  Graham's  time),  the  view  that  this  is  an  alloy  is  no 
longer  tenable,  as  this  term  is  only  strictly  applicable  to  the  union 
of  metals. 

The  absorption  of  hydrogen  by  palladium  is  readily  seen  by 
making  a  strip  of  palladium  foil  the  negative  electrode  in  an 
electrolytic  cell  containing  acidulated  water,  the  positive  pole 
being  of  platinum.  Oxygen  will  be  evolved  from  the  latter 
electrode,  while  for  some  time  no  gas  will  be  disengaged  from 
the  surface  of  the  palladium,  the  hydrogen  being  completely 
absorbed  by  the  metal.  During  the  absorption  of  hydrogen  the 
palladium  undergoes  an  increase  in  volume  :  Graham  observed  the 
increase  in  length  of  a  palladium  wire  to  be  equal  to  1.6  per  cent. 

*  According  to  Neumann  and  Strientz  (Zeitschrift  fiir  Analytische  Chemie, 
vol.  32),  one  volume  of  various  metals  in  a  fine  state  of  division  is  capable  of 
absorbing  the  following  amounts  of  hydrogen  : — 


Palladium,  black        .  502.35  vols. 

Platinum,  sponge       .  49.3       ,, 

Gold  ....  46.3 

Iron   ....  19.17     ,, 


Nickel        .  .  .  17. 57  vols. 

Copper^      .  .  .       4.5       ,, 

Aluminium  .  .       2.72     ,, 

Lead.         .  .  .       0.15     „ 


180  Inorganic  Chemistry 

This  change  in  volume  suffered  by  the  metal  may  be  strikingly 
demonstrated  by  employing  two  strips  of  palladium  foil,  protected 
on  one  side  by  a  varnish,  as  the  electrodes  in  the  electrolytic  cell. 
On  passing  the  current  the  negative  electrode  immediately  begins 
to  bend  over  towards  the  varnished  side  :  when  the  current  is 
reversed  it  again  uncurls  ;  and  the  other,  being  now  the  negative 
pole,  at  once  begins  to  perform  the  same  curling  movements. 

Hydrogen,  which  is  thus  occluded  in  the  metal  palladium,  is 
capable  of  bringing  about  a  number  of  chemical  changes  which 
ordinary  hydrogen  is  unable  to  effect  :  thus,  when  a  strip  of 
hydrogenised  palladium  is  immersed  in  a  solution  of  a  ferric  salt, 
a  portion  of  the  iron  is  reduced  to  the  ferrous  state.'5*' 

*  See  "Chemical  Lecture  Experiments,"  Nos.  27,  28,  29, 


CHAPTER   II 
OXYGEN 

Symbol,  O.     Atomic  weight  =  16.00.     Molecular  weight  =  32. 

History. — Oxygen  was  discovered  by  Priestley  (1774).  He  ob- 
tained it  by  heating  the  red  oxide  of  mercury  (known  in  those  days 
as  mercurius  calcinatus^  per  se)  by  concentrating  the  sun's  rays 
upon  it  by  means  of  a  powerful  lens.  Priestley  applied  to  the  gas 
the  name  dephlogistigated  air.  Oxygen  was  independently  dis- 
covered by  Scheele.  Scheele's  discovery  of  oxygen  was  published 
in  1775,  but  recent  research  among  his  original  papers  has  brought 
to  light  the  fact  that  the  discovery  was  actually  made  in  1773,  prior 
therefore  to  Priestley's  discovery.  Scheele  called  the  gas  empyreal 
air,  on  account  of  its  property  of  supporting  combustion.  Lavoisier 
subsequently  applied  to  this  gas  the  name  "  oxygene "  (from  ogvs, 
sour  ;  and  yevvda),  I  produce),  to  denote  the  fact  that  in  many 
instances  the  products  obtained  by  the  combustion  of  substances 
in  the  gas  were  endowed  with  acid  properties.  Oxygen,  indeed, 
came  to  be  regarded  as  an  essential  constituent  of  acids,  and  was 
looked  upon  as  the  "  acidifying  principle."  The  subsequent  deve- 
lopment of  the  science  has  shown  that  this  idea  is  erroneous,  and 
that  oxygen  is  not  a  necessary  constituent  of  all  acids. 

Occurrence. — In  the  free  state  oxygen  occurs  in  the  atmos- 
phere, mechanically  mixed  with  about  four  times  its  volume  of 
nitrogen.  In  combination  with  other  elements  it  is  found  in 
enormous  quantities.  Thus  it  constitutes  eight-ninths  by  weight 
of  water,  and  nearly  one-half  by  weight  of  the  rocks  of  which  the 
earth's  crust  is  mainly  composed. 

The  following  table  (Bunsen)  gives  the  average  composition  of 
the  earth's  solid  crust,  so  far  as  it  has  been  penetrated  by  man. 
It  must  be  remembered,  however,  that  the  greatest  depth  to  which 
man  has  examined,  when  compared  with  the  diameter  of  the  earth, 

is  after  all  only,  as  it  were,  a  mere  scratch. 

181 


1  82  Inorganic  Chemistry 

Average  Composition  of  the  EartHs  Crust. 

Oxygen  ......  44.0  to  48.7 

Silicon    ......  22.8  „  36.2 

Aluminium     .....  9.9  „  6.1 

Iron        ......  9.9  „  2.4 

Calcium  ....-,.  6.6  ,,  0.9 

Magnesium    .         .         ..         ,  2.7  „  o.i 

Sodium  .         .         .         ,         ,  2.4  „  2.5 

Potassium      .         .         .         .  1.7  „  3.1 

100.00       100.00 

Modes  Of  Formation.  —  (i.)  Oxygen  may  readily  be  obtained 
by  a  slight  modification  of  Priestley's  original  method,  namely,  by 
heating  mercuric  oxide  in  a  glass  tube,  by  means  of  a  Bunsen 
flame.  The  red  oxide  of  mercury  first  darkens  in  colour,  and  is 
decomposed  by  the  action  of  the  heat  into  mercury  and  oxygen, 
thus  — 


The  evolved  oxygen  may  be  collected  over  water  in  the  pneumatic 
trough,  while  the  mercury  condenses  in  the  form  of  metallic 
globules  upon  the  cooler  parts  of  the  tube.  This  method  of 
obtaining  oxygen  is  never  employed  when  any  quantity  of  the 
gas  is  required  —  it  is  chiefly  of  historic  interest. 

(2.)  For  experimental  purposes  oxygen  is  best  prepared  from 
potassium  chlorate.  When  this  salt  is  heated  it  melts,  and  at 
about  400°  decomposes  with  brisk  effervescence  due  to  the  evolution 
of  oxygen,  while  potassium  chloride  remains  :*  — 

KC103  =  KC1  +  3O. 

If  the  potassium  chlorate  be  previously  mixed  with  about  one- 
fourth  of  its  weight  of  manganese  dioxide,  it  gives  up  the  whole  of 
its  oxygen  at  a  temperature  considerably  below  the  melting-point 
of  the  salt,  and  at  a  greatly  accelerated  rate.  When,  therefore,  the 
oxygen  is  not  required  to  be  perfectly  pure,  a  mixture  of  these  two 

*  The  mechanism  of  this  reaction  is  more  complex  than  is  represented  by 
this  equation.  It  has  been  shown  (P.  F.  Frankland)  that  during  the  decom- 
position potassium  perchlorate,  KC1O4,  is  continuously  being  formed,  and 
again  resolved  into  KC1O3  and  O.  The  extent  to  which  this  takes  place  depend- 
ing upon  the  temperature. 


Oxygen 


183 


substances  is  usually  employed.  The  mixture  may  be  conveniently 
heated  in  a  "  Florence  "  flask,  supported  in  the  position  shown  in 
the  figure,  and  gently  heated  with  a  Bunsen  flame.  The  gas  is 
washed  by  being  passed  through  water,  and  then  collected  either 
at  the  pneumatic  trough  or  in  a  gas-holder. 

The  manganese  dioxide  is  found  at  the  end  of  the  reaction  to  be 
unchanged  :  the  part  it  plays  in  the  decomposition  belongs  to  a 
class  of  phenomena  to  which  the  name  catalysis  is  applied  ;  the 
manganese  dioxide,  in  this  instance,  being  the  catalytic  agent.  It 
was  at  one  time  supposed  that  by  its  mere  presence,  itself  under- 
going no  change,  the  manganese  dioxide  enabled  the  potassium 
chlorate  to  give  up  its  oxygen  more  readily  and  at  a  lower  tempera- 
ture ;  but  the  accumulated  evidence  which  has  been  collected  by 
the  study  of  an  increasing  number  of  similar  cases  of  catalytic 
action  leads  to  the  conclusion  that  the  manganese  dioxide  is  here 


FIG.  32. 

playing  a  more  distinctly  chemical  part  in  the  reaction.  So  far  as 
is  known,  in  all  phenomena  of  this  order,  the  catalytic  agent  is  a 
substance  which  possesses  a  certain  degree  of  chemical  affinity  for 
one  of  the  constituents  of  the  body  to  be  decomposed,  and  the 
influence  of  this  attraction  is  a  necessary  factor  in  determining  the 
splitting  up  of  the  compound.  Owing,  however,  to  certain  condi- 
tions which  are  present,  such,  for  example,  as  the  particular 
temperature  at  which  the  reaction  is  conducted,  the  catalytic  agent 
is  unable  to  actually  combine  with  the  constituent  for  which  it  has 
this  affinity,  or  if  it  combines,  the  combination  it  forms  is  unable  to 
exist  and  is  instantly  resolved  again  :  hence  the  catalytic  agent 
comes  out  of  the  action  in  the  same  state  as  it  was  at  the  com- 
mencement. 

In  the  case  before  us,  it  is  now  believed*  that  the  action  of  the 
manganese  dioxide  in  facilitating  the  evolution  of  oxygen  from 
*  Sodeau,  Trans.  Chem.  Soc.,  1902,  vol.  ii.  p.  1066. 


184  Inorganic  Chemistry 

potassium  chlorate  is  due  to  the  formation  of  a  higher  oxide  of 
manganese  by  the  oxidising  action  of  the  chlorate,  which  oxide 
being  unstable  under  the  existing  conditions,  subsequently  breaks 
up  into  oxygen  and  the  original  oxide.* 

The  temperature  at  which  this  reaction  takes  place  is  below  that 
which  is  necessary  for  the  formation  of  potassium  perchlorate, 
hence  under  these  conditions  this  salt  is  not  produced. 

(3.)  When  manganese  dioxide  itself  is  heated  to  bright  redness,  it 
parts  with  one-third  of  its  oxygen  and  is  converted  into  trimanganic 
tetroxide. 

3MnO2=  Mn3O4+  O2. 

(4.)  Other  peroxides,  when  heated,  similarly  yield  a  portion  of  the 
oxygen  they  contain.  One  of  these,  namely,  barium  peroxide,  is 
now  largely  employed  for  the  preparation  of  oxygen  upon  a  manu- 
facturing scale.  This  method,  known  as  Briris  process,  from  the 
name  of  the  inventor,  is  based  upon  the  fact  that  when  barium 
oxide  (BaO)  is  heated  in  contact  with  air,  it  unites  with  an  additional 
atom  of  oxygen,  forming  barium  peroxide,  thus  — 


And  that  when  this  substance  is  still  further  heated,  it  again  parts 
with  the  additional  oxygen  and  is  reconverted  into  the  monoxide  — 


BaO2  = 

The  process,  therefore,  is  only  an  indirect  method  of  obtaining 
oxygen  from  the  air,  the  same  quantity  of  barium  monoxide  being 
employed  over  and  over  again.  In  practice  it  was  found  that 
instead  of  effecting  the  two  reactions  by  altering  the  temperature, 
which  involved  loss  of  time  and  considerable  expense,  the  same 
result  could  be  obtained  by  altering  the  pressure  and  keeping  the 
temperature  constant.  If  the  monoxide  be  heated  to  the  lower 
temperature,  at  which  the  first  reaction  takes  place,  and  air  be 
passed  over  it  at  the  ordinary  atmospheric  pressure,  atmospheric 
oxygen  is  taken  up  and  barium  peroxide  is  formed.  If  the  pressure 

*  Secondary  reactions  simultaneously  take  place,  resulting  in  the  formation 
of  small  quantities  of  potassium  permanganate,  and  the  evolution  of  traces  of 
chlorine. 


Oxygen 


185 


be  then  slightly  reduced  by  suitable  exhaust  pumps,  the  peroxide 
immediately  gives  up  one  atom  of  oxygen  without  any  further 
application  of  heat,  and  is  retransformed  into  the  monoxide.  In 


this  way,  by  alternately  sending  air  through  the  heated  retorts 
containing  the  oxide  and  then  exhausting  the  retorts,  a  continuous 
process  is  obtained  without  change  of  temperature. 

The  modus  operandi  of  the  process  will  be  seen  from  Fig.  33, 


1 86  Inorganic  Chemistry 

which  represents  the  general  arrangement  of  the  apparatus.  A 
number  of  retorts,  R,  consisting  of  long  narrow  iron  pipes,  are 
arranged  vertically  in  rows  in  the  furnace,  where  they  are  heated 
by  means  of  "producer-gas"  (i.e.  carbon  monoxide  with  atmospheric 
nitrogen,  obtained  by  the  regulated  combustion  of  coke). 

By  means  of  the  pump  P,  air  is  drawn  in  at  the  "  air  intake,"  and 
forced  through  purifiers  in  order  to  withdraw  atmospheric  carbon 
dioxide  ;  the  complete  removal  of  this  impurity  being  essential  to 
the  successful  carrying  out  of  the  operation.  The  purifiers  are 
so  arranged  that  any  of  them  can  be  thrown  out  of  use  at 
will. 

By  means  of  automatic  gear  the  purified  air  is  sent  through  pipe 
J  to  the  distributing  valve  X,  from  which  it  passes  by  the  pipe  Y  into 
the  retorts,  being  made  to  pass  down  through  one  row  and  up  through 
the  other.  The  oxygen  is  then  absorbed,  and  the  accumulating 
nitrogen  escapes  by  the  relief  valve  W.  When  the  absorption  of 
oxygen  by  the  barium  monoxide  in  the  retorts  has  continued  for 
ten  or  fifteen  minutes,  the  automatic  reversing  gear  comes  into 
operation.  The  relief  valve  W  is  thereby  closed,  communication 
with  the  purifiers  is  cut  off,  and  the  action  of  the  pumps  at  once 
causes  a  reduction  of  pressure  within  the  retorts.  When  the  pres- 
sure falls  to  about  660  mm.  (26  inches,  or  about  13  Ibs.  on  the 
square  inch),  the  peroxide  gives  up  oxygen,  and  is  reduced  to  the 
monoxide.  The  oxygen  is  drawn  away  by  the  pipe  J  and  is  passed 
on  to  a  gas-holder.  The  first  portions  of  gas  that  are  drawn  out 
of  the  retorts  will  obviously  be  mixed  with  the  atmospheric  nitrogen 
which  was  there  present ;  in  order  that  this  shall  be  got  rid  of,  the 
automatic  gear  is  so  arranged  that  communication  with  the  pipe 
leading  to  the  gas-holder  is  not  opened  until  a  few  seconds  after 
the  reversing  gear  is  in  operation,  and  the  first  portions  of  gas  that 
are  pumped  out  are  made  to  escape  into  the  air  by  a  snifting  valve 
S,  which  is  automatically  opened  and  closed. 

(5.)  Oxygen  may  be  obtained  by  heating  manganese  dioxide  with 
sulphuric  acid,  the  dioxide  parting  with  the  half  of  its  oxygen,  and 
a  sulphate  of  the  lower  oxide  being  formed — * 

MnO2  +  H2SO4=  MnSO4  +  H2O  +  O. 

(6.)  Similarly,  potassium  dichromate  (a  salt  containing  chromium 

*  In  order  to  avoid  unnecessarily  complicating  chemical  equations,  it  is  some- 
times convenient  to  represent  them  atomically.  Moreover,  by  so  doing  the 
mechanism  of  the  reaction  is  often  rendered  more  clear. 


Oxygen  187 

trioxide,  CrO3),  when  heated  with  sulphuric  acid,  yields  oxygen  ;  the 
chromium  at  the  same  time  being  reduced  to  a  lower  state  of  oxi- 
dation, viz.,  Cr2O3,  in  which  condition  it  unites  with  sulphuric  acid, 
forming  chromium  sulphate  — 

K2Cr2Or  +  4H2SO4=K2SO4+Cr2(SO4)3  +  4H 


During  the  reaction  the  red  colour  of  the  dichromate  changes  to 
the  deep  olive-green  colour  possessed  by  chromium  sulphate. 

(7.)  Many  other  highly  oxidised  salts  yield  oxygen  when  acted 
upon  by  sulphuric  acid  ;  thus,  with  potassium  permanganate  the 
following  action  takes  place  :  — 


K2Mn2O8  +  3H2SO4=K2SO4 

(8.)  If  hydrogen  peroxide  be  added  to  dilute  sulphuric  acid,  and 
the  mixture  dropped  upon  a  solution  of  potassium  permanganate 
contained  in  a  suitable  generating  flask,  a  rapid  evolution  of  oxygen 
takes  place  at  the  ordinary  temperature,  thus  — 

K2Mn2O8  +  3H2SO4+5H2O2=K2SO4+2MnSO4+8H2O  +  5O2. 

(9.)  When  strong  sulphuric  acid  is  dropped  upon  fragments  of 
brick  or  pumice-stone,  contained  in  an  earthenware  or  platinum 
retort  and  maintained  at  a  bright  red  heat,  the  acid  is  decomposed 
into  water,  sulphur  dioxide,  and  oxygen  — 

H2SO4  =  H2O  +  SO2+O. 

The  products  of  the  decomposition  are  passed  through  water,  which 
absorbs  the  sulphur  dioxide,  and  also  arrests  any  undecomposed 
sulphuric  acid,  and  the  oxygen  is  collected  over  water.  When  this 
process  is  used  on  a  large  scale,  the  sulphur  dioxide  is  absorbed  by 
being  passed  through  a  tower  filled  with  coke,  and  down  which  a 
stream  of  water  is  allowed  to  trickle,  and  the  solution  so  obtained 
can  be  utilised  in  the  manufacture  of  sulphuric  acid. 

(10.)  Oxygen  can  be  obtained  from  bleaching-powder  by  methods 
which  afford  interesting  instances  of  catalytic  action.*  The 
composition  of  bleaching-powder  is  expressed  by  the  formula 
Ca(OCl)Cl.  If  this  substance  be  mixed  with  water,  and  a  small 
quantity  of  precipitated  cobalt  oxide  added,  and  the  mixture  gently 
warmed,  oxygen  is  rapidly  evolved.  The  cobalt  oxide,  CoO,  is  the 
catalytic  agent  ;  it  is  able  to  combine  with  more  oxygen  to  form 

*  Experiments  35,  36,  37.  152,  "  Chemical  Lecture  Experiments,"  new  ed. 


1  88  Inorganic  Chemistry 

Co2O3,  but  this  compound  is  reduced  as  fast  as  it  is  formed,  and 
the  oxygen  is  evolved  as  gas  — 

(1)  Ca(OCl)Cl  +  2CoO  =  Co2O3+CaCl2. 

(2)  Co2O3=2CoO  +  O. 

A  solution  of  calcium  hypochlorite,  which  may  be  obtained  from 
bleaching-powder  (see  Bleaching-powder),  behaves  in  the  same 
way  ;  and,  as  in  the  above  reaction,  nickel  oxide  may  be  substi- 
tuted for  cobalt  — 

Ca(OCl)2=CaCl2 


(n.)  A  similar  instance  of  catalysis,  by  which  oxygen  may  be 
obtained,  is  seen  when  a  stream  of  chlorine  gas  is  passed  through 
boiling  milk  of  lime,  to  which  a  small  quantity  of  the  oxide  of 
cobalt  or  nickel  has  been  added  — 

CaH2O2  +  C12  =  CaCl2+  H2O  +  O. 

A  reaction  of  the  same  order  takes  place  when  the  milk  of  lime  is 
replaced  by  either  potassium  or  sodium  hydroxide  — 


(12.)  When  a  mixture  of  steam  and  chlorine  gas  is  heated  to  bright 
redness,  the  steam  is  decomposed,  the  hydrogen  combining  with 
the  chlorine  to  form  hydrogen  chloride  (hydrochloric  acid),  and 
the  oxygen  is  set  free  — 


In  order  to  prepare  oxygen  by  this  reaction,  chlorine  gas  is  caused 
to  bubble  through  water  which  is  briskly  boiling  in  a  glass  flask, 
F  (Fig.  34).  The  mixture  of  chlorine  and  steam  is  then  passed 
through  a  porcelain  tube  filled  with  fragments  of  porcelain,  and 
maintained  at  a  bright  red  heat  in  a  furnace.  The  issuing  gases 
are  passed  through  a  Woulfs  bottle,  containing  a  solution  of 
sodium  hydroxide,  in  order  to  absorb  the  hydrochloric  acid,  and 
the  oxygen  is  collected  at  the  pneumatic  trough. 

(13.)  Oxygen  is  formed  on  a  large  scale  in  nature  by  the  decom- 
position of  atmospheric  carbon  dioxide  by  the  green  leaves  of 
plants,  under  the  influence  of  light.  The  carbon  dioxide  is  decom- 


Oxygen 


189 


posed  into  carbon,  which  is  assimilated  by  the  plant,  and  into 
oxygen,  which  is  thrown  into  the  atmosphere.  It  has  been  esti- 
mated that  i  square  metre  of  green  leaf  is  able,  under  the  influ- 
ence of  sunlight,  to  decompose  more  than  I  litre  of  carbon  dioxide 
per  hour. 

(14.)  Of  the  many  other  methods  by  which  it  has  been  proposed, 
from  time  to  time,  to  manufacture  oxygen  on  a  large  scale,  may  be 
mentioned  one,  known  as  the  Tessie  du  Motay  process,  from  the 
name  of  the  inventor.  This  method  consists  in  the  alternate  for- 
mation and  decomposition  of  sodium  manganate.  The  process 
consists  of  two  operations,  which  are  carried  out  at  different  tem- 
peratures. When  a  current  of  air  is  passed  over  a  moderately 


FIG.  34. 


heated  mixture   of  manganese   dioxide   and    sodium  hydroxide, 
sodium  manganate  is  formed  — 


And  if  this  sodium  manganate  be  heated  to  bright  redness,  and  a 
current  of  steam  at  the  same  time  passed  over  it,  the  manganate  is 
reduced  to  dimanganic  trioxide,  sodium  hydroxide  is  reformed,  and 
oxygen  evolved,  thus  — 


On  again  passing  air  over  the  residue,  after  allowing  the  tempera- 


190  Inorganic  Chemistry 

lure  of  the  mass  to  fall  to  that  at  which  the  first  reaction  was 
conducted,  sodium  manganate  is  once  more  reformed  — 


Properties.  —  Oxygen  is  a  colourless  gas,  having  no  taste  or 
smell.  It  is  slightly  heavier  than  air,  its  specific  gravity  being 
1.1056  (air=  i).  One  litre  of  the  gas,  at  the  standard  temperature 
and  pressure,  weighs  1.429  grammes.  Oxygen  is  slightly  soluble 
in  water.  I  c.c.  of  water  at  o°  C.  dissolves  0.0489  c.c.  of  oxygen 
measured  at  o°  C.  and  760  mm.  pressure.  The  solubility  of  oxygen 
in  water  diminishes  as  the  temperature  rises  in  accordance  with 
the  interpolation  formula  (Winkler)  :  — 

£•=0.0489  -  .001  34  1  3/4-  .0000283/2  —  .000000295  34/3. 

Fish  are  dependent  upon  the  dissolved  oxygen  in  water  for  their 
supply  of  this  gas  for  respiration.  Oxygen  is  also  soluble  in  molten 
silver,  which  is  capable  of  absorbing  about  twenty  times  its  own 
volume  of  this  gas  (see  Silver). 

Oxygen  is  endowed  with  very  powerful  chemical  affinities.  Even 
at  the  ordinary  temperature  it  is  able  to  combine  with  such  elements 
as  phosphorus,  sodium,  potassium,  and  iron.  Most  of  the  chemical 
phenomena  exhibited  by  the  atmosphere  are  due  to  the  presence  in 
it  of  free  oxygen,  the  atmosphere  being  practically  oxygen  diluted 
with  four  times  its  volume  of  nitrogen.  Thus,  when  a  piece  of  bright 
metallic  sodium  is  exposed  to  the  air,  the  surface  becomes  instantly 
tarnished  and  coated  over  with  a  film  of  oxide  :  when  iron  rusts,  it 
in  the  same  way  is  being  acted  upon  by  the  oxygen  of  the  air  forming 
an  oxide  (or  hydrated  oxide)  of  iron  ;  in  these  cases  the  metals  are 
said  to  become  oxidised.  If  the  metal  be  obtained  in  a  sufficiently 
finely  divided  condition  before  being  exposed  to  the  air,  or  to  pure 
oxygen,  this  process  of  oxidation  may  proceed  so  rapidly  that  the 
heat  developed  by  the  combination  will  cause  the  metal  to  burn. 
When  the  process  of  oxidation  is  accompanied  by  light  and  heat, 
the  phenomenon  is  known  as  combustion^  the  oxygen  being  spoken 
of  as  the  supporter  of  combustion  :  bodies  which  burn  in  the  air, 
therefore,  are  simply  undergoing  rapid  combination  with  oxygen. 
It  will  obviously  follow,  that  bodies  which  are  capable  of  burning 
in  the  air  will  burn  with  greatly  increased  rapidity  and  brilliancy 
when  their  combustion  is  carried  on  in  pure  or  undiluted  oxygen. 
A  glowing  chip  of  wood,  or  a  taper  with  a  spark  still  upon  the 


Oxygen 


191 


wick,  when  plunged  into  pure  oxygen,  will  be  instantly  rekindled. 
Such  substances  as  sulphur,  charcoal,  phosphorus,  which  readily 
burn  in  air,  when  burnt  in  pure  oxygen  carry  on  their  combustion 
with  greatly  increased  brilliancy.  Many  substances  which  are  not 
usually  regarded  as  combustible  bodies  will  burn  in  oxygen  if  their 
temperature  be  raised  sufficiently  high  to  initiate  the  combustion  ; 
thus  a  steel  watch-spring,  or  a  bundle  of  steel  wires,  if  strongly 
heated  at  one  end,  will  burn  in  oxygen,  throwing  out  brilliant 
scintillations.  This  experiment  is  most  readily  shown  by  project- 
ing a  spirit-lamp  flame  upon  the  ends  of  a  bundle  of  steel  /wire,  by 
means  of  a  stream  of  oxygen,  as  shown  in  Fig.  35.  As  soon  as  the 
ends  of  the  wire  are  sufficiently  heated,  and  begin  to  burn,  the 
lamp  may  be  withdrawn  and  the  wire  held  in  the  issuing  stream 


FIG.  35. 

of  oxygen,  in  which  it  will  continue  its  combustion  with  great 
brilliancy.* 

It  is  a  remarkable  fact,  and  one  which  has  not  yet  received  any 
entirely  satisfactory  explanation,  that  these  instances  of  combustion 
in  oxygen  will  not  take  place  if  both  the  gas  and  the  material  be 
absolutely  dry.  It  has  been  shown  that  phosphorus,  sealed  up  in  a 
tube  with  oxygen  which  has  been  absolutely  freed  from  aqueous 
vapour,  may  even  be  distilled  in  the  gas  without  any  combination 
taking  place.  A  mixture  of  oxygen  and  carbon  monoxide,  which 
under  ordinary  circumstances  explodes  when  heated,  is  found  when 
perfectly  dry  to  remain  unaffected  by  the  passage  of  an  electric 
spark  (p.  298).  Similarly,  perfectly  dry  chlorine  is  without  action 
upon  metals  such  as  copper  or  sodium,  while  under  common  con- 
ditions it  combines  with  them  with  the  greatest  readiness.  In 
all  these  cases  where  the  absolutely  dry  materials  are  incapable 

*  See  also  Experiments  48-52,  "  Chemical  Lecture  Experiments." 


1 92  Inorganic  Chemistry 

of  acting  upon  each  other,  the  introduction  of  the  minutest  trace 
of  moisture  is  sufficient  to  allow  the  action  to  proceed,  but  the 
exact  way  in  which  this  operates  in  causing  the  effect  is  at  present 
not  known  with  certainty.* 

Oxygen  is  the  only  gas  which  is  capable  of  supporting  respira- 
tion :  an  animal  placed  in  any  gas  or  gaseous  mixture  containing 
no  free  oxygen  rapidly  dies.  Undiluted  oxygen  may  be  breathed 
with  impunity  for  a  short  time,  but  its  continued  inhalation  soon 
produces  febrile  symptoms.  The  inhalation  of  oxygen  is  occa- 
sionally had  recourse  to  in  cases  of  asphyxiation,  or  under 

*  It  would  seem  evident  if  it  is  the  presence  of  water,  and  water  only  as  the 
third  substance,  which  is  the  necessary  condition  to  bring  about  chemical 
action  in  such  cases  as  these,  that  however  completely  a  mixture  of  oxygen 
and  hydrogen  were  dried,  it  would  explode  when  heated  above  the  temperature 
at  which  union  begins ;  because  the  product  of  the  combination  of  a  minute 
portion  of  the  mixture  would  furnish  sufficient  water  to  determine  the  ex- 
plosion of  the  remainder.  The  researches  of  V.  Meyer,  Dixon,  and  Baker, 
extending  over  the  last  decade,  seemed  to  entirely  confirm  this  view,  for  they 
could  detect  no  diminution  in  the  velocity  of  the  union  of  these  gases  even 
when  most  carefully  dried. 

Quite  recently,  however  (April  1902),  Baker  has  shown  that  if  these  gases 
be  perfectly  pure  (see  p.  208),  as  well  as  perfectly  dry,  the  mixture  may  be 
heated  to  temperatures  much  higher  than  that  at  which  combination  usually 
takes  place  without  exploding.  Moreover,  if  the  pure  gases  are  heated  before 
the  drying  operation  has  been  carried  to  its  highest  degree,  it  is  found  that 
although  union  begins  to  take  place  and  water  is  actually  formed  in  quantity 
greatly  in  excess  of  that  which  would  be  necessary  to  bring  about  the  action 
had  the  gases  not  been  pure,  nevertheless  no  explosion  of  the  mixture  takes 
place.  The  gases  being  absolutely  pure  to  start  with,  the  water  produced  bj» 
their  union  will  also  be  pure ;  and  it  would  appear  from  these  experiments  that 
perfectly  pure  water  alone  is  not  capable  of  bringing  about  these  chemical 
combinations. 

These  new  and  most  interesting  results  lend  support  to  the  hypothesis  which 
has  been  put  forward  in  order  to  explain  the  influence  of  water  in  determining 
such  chemical  actions  as  these,  namely,  that  chemical  action  cannot  take  place 
without  the  presence  of  an  electrolyte  ;  that  the  removal  of  water  is  in  reality 
the  removal  of  any  possibility  of  an  electrolyte  being  present.  If  the  water 
present  is  absolutely  pure,  since  pure  water  is  a  non-electrolyte  (p.  109),  it 
should  therefore  not  be  able  to  operate  in  causing  the  chemical  action  to  take 
place.  On  the  other  hand,  any  impurity  present  in  the  water  would  at  once 
cause  it  to  become  an  electrolyte,  in  which  case  it  would  be  able  to  bring 
about  the  chemical  action. 

Experiments  have  been  made  with  a  view  to  determine  whether  or  not  the 
gases  themselves  undergo  any  dissociation  in  the  moist  condition  or  during  the 
process  of  drying.  But  the  results  so  far  only  show  that  if  any  dissociation 
takes  place,  the  extent  to  which  it  occurs  is  beyond  the  limits  of  measurement 
by  any  volumetric  methods. 


Oxygen  193 

circumstances  of  great  bodily  prostration,  where  the  necessary 
oxygenation  of  the  blood  cannot  take  place  on  account  of  the 
enfeebled  action  of  the  lungs. 

Compressed  oxygen  acts  upon  the  animal  economy  as  a  poison  : 
an  animal  placed  in  oxygen  gas  under  a  pressure  of  only  a  fe\v 
atmospheres  quickly  dies. 

During  the  respiration  of  man,  air  is  drawn  into  the  lungs,  and 
is  there  deprived  of  4  to  5  per  cent,  of  its  oxygen,  and  gains  3  to  4 
per  cent,  of  carbon  dioxide.  The  oxygen  that  is  withdrawn  from 
the  inhaled  air  by  means  of  the  lungs  is  absorbed  by  the  blood. 
The  power  to  absorb  this  oxygen  is  believed  to  reside  in  a  crystal- 
line substance  contained  in  the  corpuscles  of  the  blood,  called 
hemoglobin,  with  which  it  enters  into  feeble  chemical  union,  form- 
ing the  substance  known  as  oxyhcemoglobin.  This  substance  is 
red,  and  imparts  to  arterial  blood  its  well-known  colour.  During 
its  circulation  in  the  system  the  oxyhasmoglobin  parts  with  its 
oxygen,  and  is  reconverted  into  the  purple-coloured  haemoglobin. 
Under  normal  conditions  the  whole  of  the  oxyhaemoglobin  is  not 
so  reduced,  for  venous  blood  is  found  still  to  contain  it  to  some 
extent.  The  amount  of  carbon  dioxide  exhaled  is  diminished 
during  sleep,  and  to  a  still  greater  extent  during  hibernation. 

Oxygen  can  be  liquefied  at  very  low  temperatures  by  the  appli- 
cation of  moderate  pressure  (see  Liquefaction  of  Gases).  It  was 
first  liquefied  in  1877  by  Cailletet,  and  independently  by  Pictet. 
Its  critical  temperature  is  -118.8°,  at  which  point  a  pressure  of 
58  atmospheres  is  required  to  bring  about  its  liquefaction. 

Liquid  oxygen  is  a  pale  steel-blue,  mobile  liquid,  which  boils  at 
-182.5°.  Its  specific  gravity  at  -182.5°  *s  I<I3I5-  '^ne  liquid 
expands  when  warmed  much  more  rapidly  than  gases  do  for  the 
same  increment  of  temperature,  and  its  density  diminishes  in  pro- 
portion, thus — 

At  -  182.5°  density  =1.1 24. 
„  -139°  „  =  -877- 
„  -134°  „  =  -806. 
„  -129°  „  =  .755^ 

Liquid  oxygen  is  strongly  magnetic.  If  a  quantity  of  the  liquid 
be  placed  in  a  dish  between  the  poles  of  a  powerful  electro-magnet, 
the  liquid  will  be  drawn  up  to  the  magnet  the  instant  the  latter  is 
excited. 


N 


194  Inorganic  Chemistry 


ISOMERISM — POLYMERISM — ALLOTROPY. 

Isomerism. — It  is  frequently  found  that  two  different  compounds  have  the 
same  composition  ;  that  is,  their  molecules  are  composed  of  the  same  number 
of  die  same  atoms,  and  yet  the  substances  have  different  properties.  Such 
compounds  are  said  to  be  isomeric,  the  one  is  an  isomer  of  the  other,  and  the 
phenomenon  is  called  isomerism.  Cases  of  isomerism  are  so  numerous  among 
the  compounds  of  carbon  (i.e.  in  the  realm  of  organic  chemistry,  see  Carbon, 
p.  295),  that  it  has  been  found  convenient  to  classify  them.  The  term 
isomerism,  therefore,  is  frequently  restricted  to  cases  in  which  the  compounds 
have  the  same  percentage  composition,  the  same  molecular  weight,  and  belong 
to  the  same  chemical  type  or  class  of  substances.  Thus,  the  two  compounds 
dimethyl  benzene  and  ethyl  benzene  are  both  expressed  by  the  formula 
C8H10.  The  molecules  in  each  case  contain  8  atoms  of  carbon  and  10  atoms 
of  hydrogen,  they  therefore  have  the  same  molecular  weight  and  the  same 
percentage  composition  ;  and  as  they  both  belong  to  the  same  type  or  family, 
they  are  said  to  be  isomeric  with  each  other.  The  difference  in  the  properties 
of  these  compounds  is  due  to  a  difference  in  the  arrangement  of  the  atoms 
within  the  molecules,  and  this  difference  is  expressed  in  their  formulae  in  the 
following  manner  : — 

Dimethyl  benzene,  C6H4(CH3)2.     Ethyl  benzne,  CCH,(C2H5). 

Different  compounds  having  the  same  molecular  weight  and  the  same  per- 
centage composition,  but  which  do  not  belong  to  the  same  family  of  compounds, 
are  distinguished  as  metamers.  Thus,  the  two  compounds  acetone  and  ally]' 
alcohol  are  each  expressed  by  the  formula  C3H6O.  They  have  the  same 
molecular  weight  and  the  same  percentage  composition,  but  belong  to  two 
widely  different  types  of  compounds;  they  are  therefore  called  metameric 
compounds.  The  difference  between  them  is  again  due  to  a  difference  in 
molecular  structure,  and  they  are  distinguished  by  formulas  which  convey  this 
difference,  thus : — 

Acetone,  CO(CH3)o.     Allyl  alcohol,  C3H5(HO). 

Polymerism. — This  term  is  employed  to  denote  those  cases  in  which  dif- 
ferent compounds  belonging  to  the  same  family  have  the  same  percentage 
composition,  but  differ  in  molecular  weight ;  that  is  to  say,  their  molecules  are 
composed  of  the  same  elements,  which  are  present  in  the  same  proportion ;  but 
they  do  not  contain  the  same  actual  numbers  of  the  various  atoms,  and  therefore 
have  different  weights.  Thus,  the  compounds  ethylene  (C2H4),  propylene 
(C3H6),  butylene  (C4H8),  belong  to  the  same  family,  and  have  each  the  same 
percentage  composition,  but  they  differ  in  molecular  weight.  These  sub- 
stances are  said  to  be  polymers  of  one  another. 

Allotropy  may  be  regarded  as  a  special  case  of  polymerism.  In  its  widest 
sense  the  term  is  sometimes  used  to  denote  polymerism  in  general,  but  it  is 
usually  restricted  to  those  instances  of  polymerism  which  are  exhibited  by 
elementary  bodies  only.  Many  of  the  elements  are  capable,  under  special 


Ozone  195 

conditions,  of  assuming  such  totally  different  habits  and  properties,  that  they 
appear  to  be  entirely  different  substances.  Thus,  the  element  sulphur,  as 
usually  seen,  is  a  primrose-yellow,  opaque,  solid  substance,  extremely  brittle, 
and  readily  dissolved  by  carbon  disulphide.  Under  certain  circumstances  it 
may  be  made  to  appear  a  totally  different  thing  ;  it  is  then  a  translucent  amber- 
coloured  substance,  soft  and  elastic  like  indiarubber,  and  insoluble  in  carbon 
disulphide  ;  it  is  still  sulphur,  and  sulphur  only.  Phosphorus,  again,  as  usually 
known,  is  a  nearly  colourless,  translucent,  wax-like  solid,  which  melts  at  a 
temperature  only  slightly  above  that  of  the  hand,  and  which  takes  fire  a  few 
degrees  higher;  it  is  also  extremely  poisonous.  Under  special  influences 
phosphorus  can  be  made  to  assume  the  following  properties  : — A  dark  reddish- 
brown  powder,  resembling  chocolate,  which  may  be  heated  to  250°  without 
taking  fire,  and  which  is  non-poisonous.  The  substance  is  still  phosphorus, 
and  phosphorus  only.  This  property  possessed  by  certain  of  the  elements  of 
appearing  in  more  than  one  form,  of  assuming,  as  it  were,  an  alias,  is  called 
allotropy ;  the  more  uncommon  form  being  spoken  of  as  the  allotropic  modifica- 
tion, or  the  allotrope  of  the  other. 

From  a  study  of  the  best  known  instances  of  this  phenomenon,  it  is  believed 
that  allotropy,  in  all  cases,  is  due  to  a  difference  in  the  number  of  atoms  of  the 
element  that  are  contained  in  the  molecule.  In  the  case  of  ozone,  which  is 
the  allotrope  of  oxygen,  this  is  known  to  be  the  case.  Ordinary  oxygen 
molecules  consist  of  two  atoms,  while  the  molecule  of  ozone  is  an  aggregation 
of  three  oxygen  atoms. 

OZONE. 

Molecular  symbol,  O3.     Molecular  weight =48.     Density=24. 

History. — When  an  electrical  machine  is  in  operation  a  peculiar 
and  characteristic  smell  is  noticed  in  its  vicinity.  The  same  smell 
is  sometimes  observed  in  and  about  buildings,  or  other  objects,  when 
struck  by  lightning.  In  1785  it  was  observed  by  Van  Marum  that 
when  electric  sparks  were  passed  in  oxygen,  the  oxygen  acquired 
this  peculiar  smell.  Schonbein  (1840)  showed  that  the  oxygen 
obtained  by  the  electrolysis  of  water  also  contained  this  substance 
having  a  smell,  and  he  gave  to  it  the  name  ozone,  signifying  a  smell. 
Schonbein  made  a  careful  study  of  the  substance,  and  found  that 
it  might  be  obtained  by  various  other  methods.  The  more  recent 
work  of  Andrews,  Soret,  and  Brodie  has  brought  our  knowledge  of 
the  constitution  of  ozone  to  its  present  state. 

Occurrence. — Ozone  is  present  in  the  atmosphere  in  extremely 
small  quantities  (see  Atmospheric  Ozone). 

Modes  of  Formation. — (i.)  Mixed  with  an  excess  of  oxygen, 
ozone  is  best  obtained  by  exposing  pure  dry  oxygen  to  the  influence 
of  the  silent  electric  discharge.  This  may  be  effected  by  means  of 
the  instrument  shown  in  Fig.  36,  known  as  "  Siemens'  ozone  tube." 


196 


Inorganic  Chemistry 


It  consists  of  two  concentric  glass  tubes,  A  and  B.  Tube  A  is  coated 
upon  its  inner  surface  with  tinfoil,  which  is  brought  into  metallic 
contact  with  the  binding  screw  D,  as  shown  in  the  figure.  Tube  B 
is  coated  upon  the  outer  surface,  also  with  tinfoil,  which  is  in 
metallic  connection  with  binding  screw  C.  These  two  surfaces  of 
tinfoil  are  connected  by  means  of  their  respective  binding  screws 


FIG.  36. 

• 

to  the  terminals  of  a  Ruhmkorf  coil,  and  the  slow  stream  of  oxygen 
which  is  admitted  at  E,  and  which  passes  along  the  annular  space 
between  the  two  tubes,  is  there  exposed  to  the  action  of  the  silent 
electric  discharge.  A  small  portion  of  the  oxygen  so  passing 
becomes  converted  into  the  allotropic  modification,  and  the  mixture 


FIG.  37. 

of  oxygen  and  ozone  issues  from  the  narrow  tube  at  the  opposite 
end  of  the  apparatus.   ' 

For  general  purposes  of  illustration,  a  very  simple  arrangement 
may  be  substituted  for  the  above.  It  consists,  as  shown  in  Fig.  37, 
of  a  straight  length  of  narrow  glass  tube  having  a  piece  of  platinum 
wire  down  the  inside,  which  passes  out  through  the  walls  of  the 
tube  near  to  one  end,  and  is  there  sealed  to  the  glass.  A  second 


Ozone  197 

platinum  wire  is  coiled  round  the  outside  of  the  tube,  and  these  two- 
wires  are  connected  to  the  induction  coil.  On  passing  a  slow 
stream  of  oxygen  through  the  tube,  the  issuing  gas  will  be  found  to- 
be  highly  charged  with  ozone. 

(2.)  Ozone  is  also  formed  when  an  electric  current  is  passed 
through  water  acidulated  with  sulphuric  acid.  Thus,  in  the  ordinary 
electrolysis  of  water  the  oxygen  evolved  from  the  positive  electrode 
is  found  to  contain  ozone  in  sufficient  quantity  to  be  readily  detected,, 
both  by  its  odour  and  by  other  tests. 

(3.)  During  many  processes  of  slow  oxidation  at  ordinary  tempera- 
tures, ozone  is  formed  in  varying  quantities.  Thus,  when  phos- 
phorus is  exposed  to  the  air  an  appreciable  amount  of  ozone  is 
formed.  One  or  two  short  sticks  of  freshly  scraped  phosphorus 
are  for  this  purpose  put  into  a  stoppered  bottle  containing  air,  and 
allowed  to  remain  for  a  short  time,  when  the  air  will  be  found  to 
contain  ozone. 

(4.)  Ozone  is  also  formed  during  the  combustion  of  ether  upon 
the  surface  of  red-hot  platinum.  When  a  spiral  of  platinum  wire  is 
tvarmed  in  a  gas-flame,  and  while  hot  is  suspended  over  a  small 
quantity  of  ether  contained  in  a  beaker,  the  mixture  of  ether  vapour 
and  air  undergoes  combustion  upon  the  surface  of  the  platinum, 
which  continues  in  an  incandescent  state  so  long  as  any  ether 
remains.  During  this  process  of  combustion  a  considerable  quantity 
of  ozone  is  formed.  (See  also  Peroxide  of  Hydrogen.) 

(5.)  Ozone  is  formed  during  the  liberation  of  oxygen  in  a  number 
of  the  reactions  by  which  that  gas  is  obtained  ;  thus,  from  manga- 
nese dioxide  and  sulphuric  acid  the  oxygen  that  is  evolved  contains 
sufficient  ozone  to  answer  to  the  ordinary  test.  In  the  same  way, 
by  the  action  of  sulphuric  acid  upon  barium  peroxide  or  potassium 
permanganate,  this  allotrope  is  present  with  the  ordinary  oxygen. 
that  is  evolved. 

Properties. — As  prepared  by  any  of  the  methods  described, 
ozone  is  always  mixed  with  a  large  excess  of  unaltered  oxygen, 
probably  never  less  than  about  80  per  cent,  of  the  latter  gas  being 
present.  Even  in  this  state  of  dilution  it  has  a  strong  and  rather 
unpleasant  smell,  which  rapidly  induces  headache.  When  inhaled 
it  irritates  the  mucous  membranes,  and  is  rather  suggestive  of 
dilute  chlorine. 

Ozone  is  a  most  powerful  oxidising  substance  ;  it  attacks  and 
rapidly  destroys  organic  matter  :  on  this  account  ozonised  oxygen 
cannot  be  passed  through  the  ordinary  caoutchouc  tubes,  as  these 


198  Inorganic  Chemistry 

are  immediately  destroyed  by  it.  It  bleaches  vegetable  colours, 
and  most  metals  are  at  once  acted  upon  by  it.  Even  metals  like 
mercury,  which  are  entirely  unaltered  by  ordinary  oxygen,  are 
attacked  by  ozone.  Its  action  upon  mercury  is  so  marked  in  its 
result,  that  the  presence  of  exceedingly  small  traces  of  ozone  can  be 
detected  by  it  ;  the  mercury  is  seen  to  lose  its  condition  of  perfect 
liquidity,  and  adheres  to  the  surface  of  the  glass  vessel  containing 
it,  leaving  "  tails  "  upon  the  glass.  Ozone  converts  lead  sulphide 
(PbS)  into  lead  sulphate  (PbSO4),  and  liberates  iodine  from  potas- 
sium iodide  — 


This  property  is  generally  made  use  of  for  detecting  the  presence 
of  ozone,  advantage  being  taken  of  the  fact  that  iodine,  when  set 
free  from  combination  in  the  presence  of  starch,  gives  rise  to  a 
deep  blue-coloured  compound,  the  reaction  being  one  of  extreme 
delicacy.  In  order  to  apply  this  test  for  ozone,  strips  of  paper  are 
dipped  in  an  emulsion  of  starch  to  which  a  small  quantity  of  potas- 
sium iodide  has  been  added.  These  papers  may  be  dried  and 
preserved,  and  are  usually  spoken  of  as  ozone  test  papers.  When 
one  of  these  papers  is  moistened  with  water,  and  placed  in  air 
containing  ozone,  the  iodine  is  liberated  from  the  potassium  iodide, 
and  being  in  the  presence  of  starch,  the  paper  instantly  becomes 
blue  by  the  formation  of  the  coloured  compound  of  starch.  It  will 
be  obvious  that  this  method  of  testing  for  ozone  can  only  be  relied 
upon-  when  there  is  no  other  substance  present  which  is  able  to 
decompose  potassium  iodide  ;  for  example,  when  testing  for  ozone 
in  the  atmosphere,  the  presence  of  oxides  of  nitrogen  or  peroxide 
of  hydrogen  (both  of  which  are  capable  of  liberating  iodine,  and 
are  liable  to  be  present  in  the  air)  would  materially  vitiate  the 
result  (see  also  Atmospheric  Ozone).  The  above  decomposition 
of  potassium  iodide  by  ozone  may  be  made  use  of  as  a  test  for 
ozone  in  another  way,  which,  although  less  delicate,  is  also  less 
likely  to  be  vitiated  by  the  presence  of  other  substances.  Blue 
litmus  papers  are  dipped  into  water  which  has  been  rendered  very 
feebly  acid,  and  to  which  a  small  quantity  of  potassium  iodide  has 
been  added.  The  papers  may  be  dried  and  preserved.  On 
moistening  one  of  these  papers  with  water  and  exposing  it  to 
ozone  the  iodide  is  decomposed  as  in  the  former  case,  and  the 
potassium  hydroxide  which  is  formed,  being  a  powerfully  alkaline 
substance,  converts  the  colour  of  the  litmus  from  red  to  blue. 


Ozone  199 

When  heated  to  a  temperature  of  about  250°,  ozone  is  retrans- 
formed  into  ordinary  oxygen  ;  if,  therefore,  the  ozonised  gas 
obtained  by  means  of  the  Siemens'  ozone  tube  be  passed  through 
a  glass  tube  heated  by  means  of  a  Bunsen  flame,  the  whole  of  the 
ozone  will  be  decomposed,  and  the  issuing  gas  will  therefore  be 
found  to  be  without  action  upon  the  ozone  test  papers. 

Ozone  is  also  decomposed  by  certain  metallic  oxides,  such  as 
those  of  manganese,  copper,  and  silver.  The  action  appears  to  be 
one  of  alternate  reduction  and  oxidation,  the  metallic  oxide  remain- 
ing unaltered  at  the  conclusion,  thus  — 

Ag20  +  03  =  Ag2+202. 
Ag2  +  O3  =  Ag2O  +  O2. 

The  oxidising  power  of  ozone  is  due  to  the  instability  of  the  mole- 
cule and  the  readiness  with  which  it  loses  an  atom  of  oxygen, 
leaving  a  molecule  of  ordinary  oxygen,  thus  — 


The  oxygen  molecule  is  comparatively  inert,  but  the  liberated  atom 
in  its  nascent  state  is  endowed  with  great  chemical  activity.  No 
change  of  volume  accompanies  these  processes  of  oxidation  by 
ozone,  as  the  volume  of  the  oxygen  molecule  (O2)  is  the  same  as 
that  of  the  ozone  molecule  (O3),  the  third  atom  of  oxygen  being  that 
which  enters  into  new  combination  with  the  oxidised  substance. 

Ozone  is  soluble  to  a  slight  extent  in  water,  imparting  to  the 
solution  its  own  peculiar  smell.  1000  c.c.  of  water  dissolve  about 
4.5  c.c.  of  ozone. 

Under  the  influence  of  extreme  cold,  ozone  condenses  to  liquid 
having  an  intense  blue  colour.  So  deep  is  the  colour,  that  a  layer 
of  it  2  mm.  in  thickness  is  opaque.  This  liquid  is  obtained  by 
passing  ozonised  oxygen  through  a  tube  which  is  cooled  by  being 
immersed  in  boiling  liquid  oxygen,  which  has  a  temperature  of 
—  182.5°.  At  this  temperature  the  ozone  liquefies,  but  most  of  the 
oxygen  with  which  it  was  mixed  passes  on.  In  a  higher  state  of 
purity  it  has  been  more  recently  obtained  by  first  liquefying  ozonised 
oxygen,  and  then  separating  the  more  volatile  oxygen  by  fractional 
distillation.  Liquid  ozone  boils  at  —119°.  It  is  described  by 
Olszewski  and  Dewar  as  an  extremely  explosive  substance. 

Constitution  of  Ozone.  —  The  fundamental  difference  between 
ordinary  oxygen  and  its  allotrope  ozone  lies  in  the  fact  that  the 
molecule  of  the  latter  contains  three  atoms,  while  that  of  ordinary 


2OO 


Inorganic  Chemistry 


oxygen  consists  of  only  two.  Ozone,  therefore,  is  a  polymer  of 
oxygen  ;  its  molecule  is  more  condensed,  three  atoms  occupying 
two  unit  volumes.  This  conclusion  as  to  the  constitution  of  ozone 
has  been  arrived  at  from  the  consideration  of  a  number  of  experi- 
mental facts. 

(i.)  When  oxygen  is  subjected  to  the  action  of  the  electric  dis- 
charge, it  is  found  to  undergo  a  diminution  in  volume.*  This  was 
shown  by  Andrews  and  Tait  by  means  of  the  tube  seen  in  Fig.  38. 
The  tube  was  filled  with  dry  oxygen,  which  was  prevented  from 
escaping  by  means  of  the  sulphuric  acid  contained  in  the  bent  por- 
tion of  the  narrow  tube,  which  served  as  a  manometer.  When  the 
silent  discharge  was  passed  through  the  oxygen,  a 
contraction  in  the  volume  took  place,  indicated  by 
a  disturbance  of  the  level  of  the  acid  in  the  syphon. 
When  the  tube  was  afterwards  heated  to  about 
300°  C.  and  allowed  to  cool,  the  gas  was  found  to 
have  returned  to  its  original  volume,  and  to  be 
devoid  of  ozone.  This  could  be  repeated  inde- 
finitely, the  gas  contracting  when  ozonised  and  re- 
expanding  when  the  ozone  was  converted  by  heat 
into  ordinary  oxygen.  As  only  a  very  small  propor- 
tion of  the  oxygen  was  converted  into  ozone,  this 
experiment  alone  afforded  no  clue  as  to  the  rela- 
tion between  the  change  of  volume  and  the  extent 
to  which  this  conversion  took  place. 

(2.)  A  small  sealed  glass  bulb,  containing  a  solu- 
tion of  potassium  iodide,  was  placed  in  the  tube 
before  the  experiment.  The  oxygen  was  ozonised,. 
and  the  usual  contraction  noticed.  The  bulb  was 
then  broken,  and  on  coming  in  contact  with  the  ozone  present 
the  potassium  iodide  was  decomposed,  iodine  being  liberated. 
No  further  contraction,  however,  followed  ;  and,  further,  when  the 
tube  was  subsequently  heated  to  300°  and  cooled,  the  gas  suffered 
no  increase  in  volume.  By  carefully  estimating  the  amount  of 
iodine  that  was  liberated  by  the  ozone,  the  actual  amount  of  oxygen 
which  had  caused  this  liberation  could  be  determined  according 
to  the  equation  — 


FIG.  38. 


and  it  was  found  that  the  volume  of  oxygen  so  used  up  was  exactly 
*  "  Chemical  Lecture  Experiments,"  new  ed.,  Nos.  63,  64. 


Ozone 


20 1 


equal  to  the  contraction  which  first  resulted  on  the  ozonisation  of 
the  oxygen. 

These  facts  proved  that  when  potassium  iodide  was  oxidised  by 
ozone  a  certain  volume  of  ordinary  oxygen  was  liberated,  which 
was  equal  to  the  volume  of  ozone  ; 
and  a  certain  volume  was  used  up, 
which  was  equal  to  the  original 
contraction. 

These  facts  were  explained  by 
the  supposition  that  ozone  was  repre- 
sented by  the  molecular  formula  O3  ; 
and  its  action  upon  potassium  iodide 
may  be  expressed  as  follows — 

2KI  +  H2O-f-O3  =  O2  +  l2  +  2KHO. 

(3.)  To  prove  the  correctness  of 
this  supposition,  however,  it  was 
necessary  to  learn  the  exact  relation 
between  these  two  volumes.  This 
Soret  did,  by  making  use  of  the  pro- 
perty possessed  by  turpentine  (and 
other  essential  oils)  of  absorbing 
ozone  without  decomposing  it  ;  and 
he  found  that  the  diminution  in 
volume  which  took  place  by  absorb- 
ing ozone  from  ozonised  oxygen  wras 
exactly  twice  as  great  as  the  increase 
in  volume  that  resulted  when  the 
same  volume  of  ozonised  oxygen 
was  heated. 

This  fact  may  be  shown  by  means 
of  the  apparatus,  Fig.  39.*  The 
oxygen  to  be  ozonised  is  contained 
in  the  annular  space  between 
the  elongated  hollow  stopper,  which 
reaches  nearly  to  the  bottom,  and 
the  outer  tube.  The  turpentine  is 
contained  in  a  little  sealed  thin  glass  tube  d,  almost  capillary  in  bore, 
which  is  held  in  position  between  four  little  projecting  glass  points  a 
and  b  upon  the  stopper  and  outer  tube.  The  temperature  is  main- 
*  Newth,  Trans.  Chem.  Soc.t  1896,  p.  1298. 


B 


FIG.  39. 


2O2  Inorganic  Chemistry 

tained  constant  throughout  the  experiment  by  placing  the  apparatus 
in  melting  ice.  One  wire  from  the  induction  coil  is  dipped  into  the 
ice  water,  while  the  other  passes  into  the  dilute  acid  contained  in  the 
stopper.  When  the  electric  discharge  is  passed  a  portion  of  the 
oxygen  is  ozonised,  resulting  in  a  contraction  in  the  volume  which 
is  indicated  by  a  rise  of  the  liquid  in  the  gauge  ;  when  sufficient 
contraction  has  taken  place  the  discharge  is  interrupted,  and  the 
contents  of  the  capillary  tube  brought  into  contact  with  the  gas. 
This  is  done  by  a  slight  twist  of  the  stopper,  which  thereby  crushes 
the  little  tube  and  throws  out  the  turpentine.  Immediately  a 
further  contraction  takes  place,  due  to  the  absorption  of  the  ozone 
by  the  reagent,  and  if  the  gauge  be  graduated  it  will  be  seen  that 
this  second  contraction  is  twice  as  great  as  the  first. 

(4.)  If  the  molecule  of  ozone  be  correctly  represented  by  O3,  its 
density  will  be  24,  as  against  16  for  oxygen  ;  and  its  rate  of  diffu- 
sion will  be  proportionately  slower  in  accordance  with  the  law  of 
gaseous  diffusion  (see  Diffusion  of  Gases,  p.  84).  Soret  found  that 
this  was  actually  the  case,  and  from  his  experiments  the  number  24 
for  the  density  of  ozone  receives  conclusive  confirmation  (see  also 
p.  228), 


CHAPTER   III 

COMPOUNDS    OF   HYDROGEN    WITH    OXYGEN 
THERE  are  two  oxides  of  hydrogen  known,  viz.  : — 

Hydrogen  monoxide,  or  water        .         .         .         .     H0O 
Hydrogen  dioxide  .......     H.,O2 

WATER. 

Formula,  K2O.     Molecular  weight  =  18.02. 

Until  the  time  of  Cavendish,  water  was  considered  to  be  an 
elementary  substance.  Priestley  had  noticed  that  when  hydrogen 
and  oxygen  were  mixed  and  inflamed,  moisture  was  produced, 
and  he  had  also  observed  that 
the  water  so  obtained  was  some- 
times acid.  Cavendish  showed 
that  the  water  was  actually  the 
product  of  the  chemical  union  of 
hydrogen  with  oxygen,  and  he 
also  discovered  that  the  acidity 
which  this  water  sometimes  pos- 
sessed was  due  to  the  presence 
of  small  quantities  of  nitric  acid  ; 
and  he  traced  the  formation  of 
this  acid  to  the  accidental  pre- 
sence of  nitrogen  (from  the  at- 
mosphere) with  which  the  gases 
were  sometimes  contaminated. 

Cavendish  filled  a  graduated 
bell-jar  with  a  mixture  of  hydro-  pIG  40 

gen  and  oxygen,  in  the  propor- 
tion of  two  volumes  of  the  former  to  one  of  oxygen,  and  he  attached 
to  the  bell-jar   a  stout  glass  vessel   resembling   the   pear-shaped 

apparatus  shown  in  Fig.  40,  which  was  perfectly  dry  and  rendered 

203 


204  Inorganic  Chemistry 

vacuous.  On  opening  the  stop-cocks,  gas  entered  the  exhausted 
tube,  which  was  furnished  at  the  top  with  two  platinum  wires.  The 
cocks  were  again  closed  and  an  electric  spark  passed  through  the 
mixed  gases,  thereby  causing  their  explosion,  when  the  interior 
surface  of  the  previously  dry  glass  vessel  was  found  to  be  dimmed 
with  a  film  of  moisture.  On  again  opening  the  stop-cocks  more 
gas  was  drawn  into  the  upper  vessel,  the  same  volume  passing  in 
as  originally  entered  the  evacuated  apparatus.  This  showed  that 
the  two  gases  in  their  combination  with  each  other  had  entirely 
disappeared.  By  repeatedly  filling  the  vessel  with  the  mixed  gases 
and  causing  them  to  unite  in  this  way,  Cavendish  succeeded  in 
collecting  sufficient  of  the  water  to  identify  the  liquid,  and  prove 
that  it  was  in  reality  pure  water. 

The  more  exact  volumetric  proportion  in  which  oxygen  and 
hydrogen  combine  to  form  water  has  been  determined  by  modern 
eudiometric  methods  which  have  been  developed  from  Cavendish's 
experiment.  Accurately  measured  volumes  of  the  two  gases  are 
introduced  into  a  long  graduated  glass  tube  standing  in  the 
mercurial  trough  and  provided  with  two  platinum  wires,  by  means 
of  which  an  electric  spark  can  be  passed.  The  gases  are  caused  to 
unite  by  means  of  the  spark,  and  the  contraction  in  volume  is 
carefully  observed.  Fig.  41  shows  the  apparatus  for  this  purpose. 
The  long  glass  tube  A  having  a  millimetre  scale  graduated  upon  it, 
and  having  two  platinum  wires  sealed  into  the  glass  near  the  upper 
and  closed  end,  is  completely  filled  with  mercury  and  inverted  in 
the  trough  of  the  same  liquid  :  this  tube  is  known  as  a  eudiometer. 
A  quantity  of  pure  oxygen  is  then  introduced  into  the  tube,  and 
the  volume  occupied  by  the  gas  carefully  read  off  upon  the  gradua- 
tions. Seeing  that  the  volume  occupied  by  a  given  mass  of  gas  is 
dependent  both  upon  the  temperature  and  the  pressure,  each  of 
these  factors  has  to  be  taken  into  account  in  the  process  of  this 
experiment.  The  temperature  is  ascertained  by  the  attached 
thermometer  T.  The  pressure  under  which  the  gas  is,  will  be  the 
atmospheric  pressure  at  the  time  (ascertained  by  the  barometer  B 
placed  near  the  apparatus)  minus  the  pressure  of  a  column  of 
mercury,  equal  to  the  height  of  the  mercury  within  the  eudiometer 
above  the  level  of  that  in  the  trough.  This  height  is  obtained  in 
millimetres  by  carefully  reading  upon  the  graduated  scale  the  level 
of  the  mercury  in  the  trough  and  the  top  of  the  column  in  the 
tube,  and  the  number  of  millimetres  so  obtained  is  deducted  from 
the  barometric  reading.  These  observations  are  made  by  means  of 


Water 


205 


a  telescope  placed  at  such  a  convenient  distance  that  the  heat  of 
the  body  may  not  introduce  disturbances. 

The  data  obtained  give  the  volume  of  gas  at  a  particular  tem- 
perature, and  under  a  pressure  less  than  that  of  the  atmosphere 
By  the  process  of  calculation  explained  under  the  general  pro- 
perties of  gases  (p.  69),  this  is  reduced  to  the  standard  temperature 
and  pressure,  viz.,  o°  and  760  mm. 

A  quantity  of  hydrogen  is  then  introduced  into  the  eudiometer, 
considerably  in  excess  of  that  required  for  complete  combination 


FIG.  41. 

•with  the  oxygen,  and  the  volume  again  ascertained  with  the  above 
precautions  and  corrections. 

The  difference  between  the  first  and  second  reading  will  give  the 
volume  of  hydrogen  which  has  been  added. 

The  eudiometer  is  then  firmly  held  down  against  a  pad  of  caout- 
chouc upon  the  bottom  of  the  trough,  and  the  gases  fired  by  an 
electric  spark  from  a  Ruhmkorff  coil.  A  bright  flash  of  light 
passes  down  the  tube,  and  on  releasing  it  from  the  indiarubber  bed, 
mercury  enters  to  fill  the  space  previously  occupied  by  the  gases 
which  have  combined. 


206 


Inorganic  Chemistry 


On  allowing  the  instrument  to  once  more  acquire  the  tempera- 
ture of  the  surrounding  atmosphere,  the  residual  volume  is  read  off 
and  corrected  for  temperature  and  pressure. 

The  following  data  have  now  been  obtained  : — 

(i.)  The  volume  of  oxygen,  corrected  for  temperature  and 
pressure. 

(2.)  The  volume  of  mixed  oxygen  and  hydrogen,  corrected  for 
temperature  and  pressure. 

(3.)  The  volume  of  residual  hydrogen,  corrected  for  tempera- 
ture and  pressure. 

A  concrete  example  will  explain  how  the  result  is  deduced  from 
these  observations  : — • 

Corrected  volume  of  oxygen  used 45-35 

Corrected  volume  after  the  addition  of  hydrogen     .         .     256.05 
Corrected  volume  of  residual  hydrogen   ....     120.10 

256.05—   45.35  =  210. 7o  =  total  volume  of  hydrogen  employed. 
210.70—120.10=:   9O.6o=volume  of  hydrogen  which  has  combined  with 

45-35  volumes  of  oxygen. 

.'.  45.35:1:190.60:1.997. 
.'.  One  volume  of  oxygen  has  combined  with  1.997  volume  of  hydrogen 

to  form  water.* 

The  volume  composition  of  water  may  be  shown  by  analytical 
processes,  as  well  as  the  synthetical 
method  described  above.  This  decom- 
position of  water  is  most  conveniently 
effected  by  means  of  an  electric  cur- 
rent. If  the  two  terminals  from  a  gal- 
vanic battery  are  connected  to  two 
pieces  of  platinum  wire  or  foil,  and 
these  are  dipped  into  acidulated  water, 
bubbles  of  gas  make  their  appearance 
upon  each  of  the  wires.  If  these  two 
strips  of  platinum  be  so  arranged  in  a 
bottle  that  all  the  gas  evolved  escapes 
by  a  delivery-tube  (Fig.  42),  it  will 
be  found  that  the  gas  explodes  violently  on  the  application  to  it 


FIG.  42. 


*  In  accurate  experiments  the  volume  occupied  by  the  minute  quantity  of 
water  formed  has  to  be  taken  into  account,  and  a  number  of  other  corrections 
have  to  be  made  that  are  not  mentioned  in  this  outline  description  of  the 
process. 


Water 


207 


of  a  lighted  taper,  showing  it  to  be  a  mixture  of  oxygen  and  hy- 
drogen. By  modifying  the  apparatus  in 
such  a  way  that  the  gas  from  each  platinum 
plate  shall  be  collected  in  separate  tubes, 
so  arranged  that  the  volumes  of  the  gases 
can  be  measured,  it  is  found  that  twice  as 
much  hydrogen  is  evolved,  in  a  given  time, 
as  oxygen.  A  convenient  form  of  volta- 
meter is  seen  in  Fig.  43,  where  the  two 
measuring  tubes  are  suspended  over  the 
platinum  plates  contained  in  a  glass  basin. 
The  electrode,  which  is  connected  with  the 
negative  terminal  of  the  battery,  is  the  one 
from  which  the  larger  volume  of  gas,  viz., 
the  hydrogen,  is  evolved,  while  the  oxygen 
is  liberated  at  the  positive  plate. 

When  the  volumes  of  the  gases  are  care- 
fully measured,  it  is  found  that  they  are  not 
exactly  in  the  proportion  of  two  of  hydro- 
gen to  one  of  oxygen,  but  that  the  oxygen 
is  in  deficit  of  this  proportion.  This  is  due, 
in  the  first  place,  to  the  greater  solubility 
of  oxygen  in  water  than  of  hydrogen  ;  and, 
secondly,  to  the  formation  of  a  certain 
quantity  of  ozone  during  the  electrolysis, 
whereby  there  is  a  shrinking  of  volume  in 
the  proportion  of  three  to  two.  FIG.  43. 


iff 


This  process  of  electrolysis  has  already  been  partially  explained  en  pp.  88, 

-t- 
89.     In  the  dilute  sulphuric  acid  employed  the  ions  present  are  mainly  H  and 

SO4  (see  persulphuric  acid,  p.  424).  The  H  ions  convey  the  current  to  the 
cathode,  where  they  discharge  their  electricity  and  unite  to  form  molecules  of 

ordinary  hydrogen  gas.  The  SO4  ions  travel  to  the  anode,  but  instead  of 
regarding  them  as  becoming  discharged  and  then  interacting  with  water 
molecules  to  form  H2SO4  with  liberation  of  oxygen  gas,  the  view  now  taken 
is  that  the  minute  proportion  of  dissociated  water  molecules  plays  a  part  in 
what  actually  goes  on.  In  the  neighbourhood  of  the  anode  there  are  a  few 

O  ions  due  to  this  dissociation  of  water  molecules,  and  it  is  believed  that 

these,  and  not  the  SO4  ions,  give  up  their  charges  and  escape  as  oxygen 
molecules  ;  while  as  fast  as  they  do  so  more  H2O  molecules  dissociate  in 

order  to  furnish  sufficient  H  ions  to  establish  equilibrium  with  the  SO4  ions. 


208  Inorganic  Chemistry 

Instead,  therefore,  of  the  chemical  equation 

2S04  +  2H2O=2H2S04  +  02 
•we  may  substitute  the  ionic  equation 


HH.;+oa. 

Similarly  when  "  secondary  reactions  "  take  place  at  the  cathode,  as,  for  in- 
stance, when  sodium  chloride  or  sodium  sulphate  are  electrolysed,  it  is  believed 

+  + 

that  as  water  is  also  slightly  dissociated  into  H  and  OH  ions,  it  is  these  H 
ions  which  actually  discharge  and  escape  as  hydrogen  gas,  and  as  fast  as  they 

do  so  more  water  molecules  dissociate  so  as  to  supply  OH  ions  to  establish  equi- 

librium with  the  K  ions.     The  two  conceptions  of  the  mechanism  of  the  action 
are  expressed  by  the  equations  — 

2K  +  2H20=2KHp+Ho. 

2K  +  2:  H  ,  OH!  =  2  K  ,  OH  ;  -I-  Ha. 

It  has  been  recently  shown*  that  the  purest  "electrolytic  gas," 
as  this  mixture  of  hydrogen  and  oxygen  is  called,  is  obtained  by 
the  electrolysis  of  pure  barium  hydroxide.  Under  these  circum- 
stances the  oxygen  contains  no  ozone  or  hydrogen  peroxide. 

The  Volumetric  Composition  of  Steam.  —  When  a  mixture  of 
oxygen  and  hydrogen  is  exploded  in  a  eudiometer,  we  have  seen 
that  a  certain  contraction  of  volume  follows,  due  to  the  formation 
of  water  by  the  uniting  gases.  The  oxygen  and  hydrogen  that 
have  entered  into  combination  have  disappeared  as  gases,  the 
volume  of  the  resultant  water  being  practically  negligible.  It  is 
important  to  know  what  relation  exists  between  the  volume  of 
the  uniting  gases  and  the  volume  of  the  product  of  their  combina- 
tion when  in  a  state  of  vapour  —  that  is  to  say,  what  volume  of 
steam  is  produced  by  the  union  of  one  volume  of  oxygen  with 
two  volumes  of  hydrogen  ;  in  other  words,  whether  there  is  any 
molecular  contraction  in  the  formation  of  steam. 

To  ascertain  this,  the  mixed  gases,  in  the  exact  propoitions  to 
form  water,  must  be  made  to  combine  under  such  circumstances 
that  the  product  shall  remain  in  a  state  of  gas  or  vapour,  so  that 
its  volume  and  that  of  the  mixed  gases  may  be  measured  under 
comparable  conditions.  For  this  purpose  a  mixture  of  oxygen  and 
hydrogen,  obtained  by  the  electrolysis  of  acidulated  water,  is  in- 
troduced into  the  closed  limb  of  the  U-shaped  eudiometer  shown 
in  Fig.  44.  t  This  tube  is  graduated  into  three  equal  divisions, 
indicated  by  the  broad  black  bands,  and  is  furnished  with  two 
platinum  wires  at  the  closed  end.  It  is  also  surrounded  by  an 
outer  tube,  so  that  a  stream  of  vapour  from  some  liquid,  boiling 

*  Baker,  Jour.  Chem.  Soc,,  April  1902. 

f  See  Experiments  Nos.  74  and  75,  "  Chem.  Lecture  Experiments,"  newed. 


Water  209 

above  the  boiling-point  of  water,  can  be  made  to  circulate.  The 
most  convenient  liquid  for  the  purpose  is  amyl  alcohol,  which 
boils  at  130°.  In  this  way  the  eudiometer  and  the  contained  gases 
will  be  maintained  at  a  constant  temperature  high  enough  to  keep 
the  water  formed  by  their  combination  in  the  state  of  vapour. 

The  amyl  alcohol  is  briskly  boiled  in  the  flask,  and  its  vapour  is 
led  into  the  tube  surrounding  the  eudiometer.  The  temperature  of 
the  mixed  gases  is  thereby  raised  to  130°,  and  they  occupy  the 


Fis.  44. 

three  divisions  of  the  tube  when  the  mercury  in  the  open  limb  is 
at  the  same  level,  that  is,  when  the  gases  are  under  atmospheric 
pressure.  The  amyl  alcohol  vapour  leaves  the  apparatus  by  the 
glass  tube  at  the  bottom,  and  is  conveyed  away  and  condensed. 
An  electric  spark  is  then  passed  through  the  gases  by  means  of  the 
induction  coil.  (In  order  to  prevent  the  mercury  from  being 
forcibly  ejected  from  the  open  limb  of  the  U-tube  at  the  moment  of 
explosion,  an  additional  quantity  of  mercury  is  poured  in,  and  the 
open  end  is  closed  by  the  thumb  when  the  spark  is  passed.)  On 
bringing  the  enclosed  gas  again  to  the  atmospheric  pressure,  by 

o 


2IO 


Inorganic  Chemistry 


adjusting  the  level  of  the  mercury  until  it  is  once  more  at  the  same 
height  in  each  limb,  it  will  be  found  that  the  mercury  in  the  eudio- 
meter is  now  standing  at  the  second  band  ;  that  is  to  say,  the  three 
volumes  of  gas  originally  present  have  now  become  two  volumes  of 
steam.  This  condensation  is  expressed  in  the  molecular  equation  — 


The  Gravimetric  Composition  of  Water.  —  Having  learned 
the  composition  of  water  by  volume,  and  knowing  also  that  the 
relative  weights  of  equal  volumes  of  oxygen  and  hydrogen  are  as 
15.88  :  I,  the  composition  by  weight  can  readily  be  calculated,  thus  — 

1  volume  of  oxygen        =  1  5.88 

2  volumes  of  hydrogen  =    2.00 

17.88 
17.88  parts  by  weight  of  water  are  composed  of  2.00  parts  by 


1°. 


FIG.  45. 

weight  of  hydrogen  and  15.88  parts  of  oxygen,  or,  expressed  centesi- 
mally,  we  have — 

Oxygen     ....         88.8 
Hydrogen         .        ,        .         11.2 

100.0 

The  composition  of  water  by  weight  has  been  experimentally 
determined  with  great  care  by  a  number  of  chemists. 

The  apparatus  shown  in  Fig.  45  represents  the  method  employed 
by  Dumas  (1843).  When  copper  oxide  is  heated  in  a  stream  of 


Water  2  1  1 

hydrogen,  the  copper  oxide  is  deprived  of  its  oxygen,  which  unites 
with  the  hydrogen  to  form  water  — 


Dumas'  method  is  based  upon  this  reaction.  A  weighed  quantity 
of  perfectly  dry  copper  oxide  was  heated  in  the  bulb  A,  in  a  current 
of  hydrogen  generated  from  zinc  and  sulphuric  acid  in  the  bottle  H, 
and  rendered  absolutely  pure  and  dry  by  its  passage  through  a 
series  of  tubes  containing  absorbents.  The  water  formed  by  the 
union  of  the  hydrogen  with  the  oxygen  of  the  copper  oxide  was 
collected  in  the  second  bulb,  B,  previously  weighed  ;  and  the  un- 
condensed  aqueous  vapour  which  was  carried  forward  in  the  stream 
of  hydrogen  was  arrested  in  the  weighed  tubes  which  follow.  The 
increase  in  weight  of  the  bulb  B  and  the  weighed  tubes  gave  the 
total  weight  of  water  produced  ;  while  the  loss  of  weight  suffered 
by  the  copper  oxide  gave  the  weight  of  oxygen  contained  in  that 
water.  The  difference  between  these  two  weights  is  the  weight 
of  the  hydrogen  that  entered  into  combination  with  the  oxygen. 

As  a  mean  of  many  experiments  it  was  found  that  in  the  forma- 
tion of  236.36  grammes  of  water  the  oxygen  given  up  by  the 
copper  oxide  was  210.06  grammes. 

236.36-210.06  =  26.30, 
therefore  236.36  grammes  of  water  were  made  up  of 

Hydrogen  =   26.30 
Oxygen     =  2  1  0.06 

236.36 
The  ratio  of  hydrogen  to  oxygen  is  therefore  as  2  :  15.88. 

Hydrogen  prepared  from  zinc  and  sulphuric  acid  is  liable  to  contain  traces  of 
(i.)  Hydrogen    sulphide.     This   is   absorbed   in  the   first   tube  containing 

broken  glass  moistened  with  a  solution  of  lead  nitrate. 

(2.)  Arsenic  hydride  (absorbed  in  the  second  tube,  filled  with  glass 

(3.)  Hydrogen  phosphide  {      moistened  with  silver  sulphate. 

rabsorbed  in  the  third  tube,  containing  in  one  limb 
(4.)  Sulphur  dioxide  (  pumice  moistened  with  a  solution  of  potassium 
(5.)  Carbon  dioxide  1  hydroxide,  and  in  the  other  fragments  of  solid 

I     potassium  hydroxide. 

Tubes  4,  5,  6,  and  7.  containing  solid  potassium  hydroxide  and  phosphorus 
pentoxide  (the  two  latter  being  placed  in  a  freezing-mixture),  are  for  the  pur- 
pose of  withdrawing  every  trace  of  aqueous  vapour.  Tube  8  was  weighed  before 


212 


Inorganic  Chemistry 


and  after  the  experiment  in  order  to  test  the  absolute  dryness  of  the  hydrogen 
that  entered  the  bulb.  In  order  to  get  rid  of  dissolved  air,  the  dilute  sulphuric 
acid  used  was  previously  boiled.  Tubes  9,  10,  n  were  weighed  both  before  and 
after  the  experiment ;  while  tube  12,  which  was  not  weighed,  was  placed  at  the 
end  to  prevent  any  absorption  of  atmospheric  moisture  by  the  weighed  tubes. 
Since  the  time  of  Dumas  this  subject  has  been  reinvestigated  by  other 
experimenters,  who  have  introduced  various  modifications  into  the  process ; 
thus,  with  a  view  to  finding  the  weight  of  hydrogen  directly  and  of  eliminating 
many  of  the  possible  sources  of  error  arising  from  the  presence  of  impurities  in 


Ftc.  46. 


the  hydrogen,  the  hydrogen  has  been  absorbed  by  palladium.  The  metal  so 
charged  with  hydrogen  can  be  weighed  before  and  after  the  experiment,  and 
the  actual  "weight  of  hydrogen  used  directly  ascertained. 

Most  recently  the  matter  has  been  investigated  by  Scott  and  Rayleigh,  and 
the  results  obtained  show  only  the  slightest  departure  from  the  numbers 
obtained  by  Dumas. 

Properties  of  Water. — Pure  water  is  a  tasteless  and  odourless 
liquid.  When  seen  in  moderate  quantities  it  appears  to  be  colour- 
less, but  when  viewed  through  a  stratum  of  considerable  thickness 
it  presents  a  beautiful  greenish-blue  colour.  This  colour  may  be 


Water  213 

seen  by  filling  a  horizontal  tube  about  15  feet  long  with  the  purest 
water,  and  passing  a  strong  beam  of  light  through  it.  It  may  also 
be  perceived  by  directing  a  ray  of  light  through  a  tall  cylinder  of 
water  in  the  manner  shown  in  the  figure,  andcausing  it  to  be  reflected 
up  through  the  water  from  the  surface  of  a  layer  of  mercury  at  the 
bottom  ;  the  immerging  ray,  being  then  reflected  upon  a  screen, 
shows  the  characteristic  colour  of  the  water.  By  intercepting  the 
ray  by  a  hand  mirror  at  A,  the  white  light  can  be  thrown  upon  the 
screen  as  a  contrast  to  the  greenish-blue  tint. 

Aitkin  has  recently  shown  that  the  presence  of  extremely  finely 
divided  suspended  matters  in  water  will  give  to  the  liquid  the  appear- 
ance of  a  blue  colour.  Thus,  in  tanks  where  water  is  being  softened 
by  the  addition  of  milk  of  lime,  after  the  bulk  of  the  precipitated  chalk 
has  settled,  and  only  the  finest  particles  still  remain  suspended  in 
the  liquid,  it  is  often  noticed  that  the  water  appears  to  have  a  rich 
blue  colour.  The  wonderful  blue  colour  of  the  waters  of  many  of 
the  Swiss  lakes  is  probably  due  in  part  to  this  optical  phenomenon 
as  well  as  to  the  intrinsic  colour  of  the  water.  When  a  mass  of 
pure  snow,  such  as  falls  in  high  mountainous  regions,  is  broken 
open  in  such  a  way  that  the  light  is  reflected  from  side  to  side 
of  the  small  crevice,  the  true  greenish-blue  colour  of  the  water  is 
very  manifest. 

Water  is  compressible  to  only  a  very  slight  extent ;  thus,  under 
an  additional  pressure  of  one  atmosphere,  i coo  volumes  of  water 
become  999.95  volumes. 

Small  as  this  compressibility  is,  it  exerts  an  important  influence  upon  the 
distribution  of  land  and  water  upon  the  earth.  It  has  been  calculated  that 
owing  to  this  compression,  where  the  ocean  has  a  depth  of  six  miles,  its  surface 
is  lower  by  620  feet  than  it  would  be  if  water  were  absolutely,  non-compressible; 
and,  calculated  from  the  average  depth  of  the  sea,  its  average  level  is  depressed 
116  feet.  The  effect  of  this  depression  of  the  sea-level  is  that  2,000,000  square 
miles  of  land  are  now  uncovered  which  would  otherwise  be  submerged  beneath 
the  ocean. 

Water  is  an  extremely  bad  conductor  of  heat.  A  quantity  of  water 
contained  in  a  tube  held  obliquely  may  be  boiled  by  the  application 
of  heat  to  the  upper  layers  without  appreciably  affecting  the 
temperature  of  the  water  at  the  bottom  ;  a  fragment  of  weighted 
ice  sunk  to  the  bottom  will  remain  for  a  long  time  unmelted,  while 
the  water  a  few  inches  above  it  is  vigorously  boiling.  This  low 
conductivity  for  heat  is  shared  in  common  by  all  liquids  that  are 
not  metallic.  Indeed,  Guthrie  has  shown,  that  water  conducts  heat 


214  Inorganic  Chemistry 

better  than  any  other  substance  which  is  liquid  at  the  ordinary 
temperature,  with  the  exception  of  mercury. 

Steam. — Under  a  pressure  of  760  mm.,  water  boils  at  100° 
(see  p.  128),  and  is  converted  into  a  colourless  and  invisible  gas 
or  vapour.  The  visible  effect  that  is  observed  when  steam  is 
allowed  to  issue  into  the  atmosphere  is  due  to  the  condensation  of 
the  steam  in  the  form  of  minute  drops  of  water.  What  is  popularly 
called  steam  is  in  reality,  therefore,  not  steam,  but  an  aggrega- 
tion of  small  particles  of  liquid  water.  The  invisibility  of  steam 
is  readily  demonstrated  by  boiling  a  small  quantity  of  water  in  a 
capacious  flask  ;  as  the  steam  issues  from  the  neck  it  condenses 
in  contact  with  the  cool  air  and  presents  the  familiar  appearance, 
but  within  the  flask  it  will  be  perfectly  transparent  and  invisible. 

ICO. — At  a  temperature  of  o°  water  solidifies  to  a  transparent 
crystalline  mass.  In  the  act  of  solidification  the  water  expands 
by  nearly  ^th  of  its  volume,  10  volumes  of  water  become  10.908 
volumes  of  ice  :  solid  water,  therefore,  is  specifically  lighter  than 
liquid  water,  and  floats  upon  its  surface.  Water  in  this  respect  is 
anomalous,  for  in  the  case  of  most  other  substances  the  solid  form 
is  denser  than  the  liquid.  The  disruptive  force  exerted  by  water  at 
the  moment  of  freezing  is  the  cause  of  the  bursting  of  pipes  and 
other  vessels  containing  water  during  winter ;  and  it  is  also  an 
important  factor  in  the  economy  of  nature  in  the  disintegration  of 
rocks  and  of  soil.  Under  certain  conditions  water  may  be  cooled 
many  degrees  below  o°  without  solidification  taking  place.  Thus, 
if  a  small  quantity  of  water  contained  in  a  vacuous  tube  be  care^ 
fully  cooled  without  being  subjected  to  vibration,  its  temperature 
may  be  lowered  to  —15°  without  it  solidifying  ;  a  slight  shock, 
however,  at  once  causes  it  to  pass  into  the  solid  state,  when  its 
temperature  instantly  rises  to  o°  (see  p.  137).  Although  the  exact 
temperature  at  which  water  freezes  is  liable  to  uncertainty  from 
this  cause,  the  point  at  which  ice  melts  is,  under  ordinary  cir- 
cumstances, constant,  viz.,  o°.  Under  increased  pressure  ice 
will  melt  at  temperatures  below  o°  ;  thus,  Mousson  found  that, 
under  a  pressure  of  13,000  atmospheres,  ice  melted  at  —  18°.  The 
melting-point  of  ice  is  lowered  by  about  0.0074°  by  each  additional 
atmosphere  of  pressure  (see  p.  137). 

Between  the  temperatures  of  +4°  and  100°,  water  follows  the 
ordinary  laws  that  govern  the  expansion  and  contraction  of  liquids 
due  to  change  of  temperature  ;  if  water  be  cooled  from  100°,  it 
gradually  contracts  until  the  temperature  reaches  4°.  Between 


Water  215 

this  point  and  o°  it  forms  a  remarkable  exception  to  the  general 
law,  for,  when  cooled  below  4°,  it  slowly  expands  instead  of  con- 
tracting, and  continues  expanding  until  o°  is  reached,  when  it 
solidifies.  At  4°,  therefore,  water  expands  whether  it  be  heated 
or  cooled ;  consequently,  at  this  point  it  is  denser  than  at  any 
other  temperature.  This  temperature  is  known  as  its  point  of 
maximum  density.  (The  most  accurate  observations  fix  the  exact 
point  at  3.945°.) 

The  following  table  shows  the  change  of  volume  suffered  by 
water  on  being  heated  from  o°  to  8° : — 

i. oooooo  volumes  at      o°  becomes 

0.999915  „  +2°  „ 

0.999870  „  4°         „ 

0.999900  „  6°        „ 

I. OOOOOO  „ 

One  cubic  centimetre  of  water,  measured  at  its  point  of  maxi- 
mum density  and  at  760  mm.,  is  the  unit  of  weight  of  the  metrical 
system,  and  is  called  a  gramme* 

It  is  also  at  this  temperature  that  water  is  taken  as  the  unit  for 
comparison  of  the  densities  of  other  liquids  and  of  solids  ;  thus, 
when  it  is  stated  that  the  density  or  specific  gravity  of  diamond 
is  3.5,  it  is  meant  that  diamond  is  3.5  times  as  heavy  as  an  equal 
bulk  of  water  measured  at  its  point  of  maximum  density. 

The  fact  that  water  has  a  point  of  maximum  density  remote  from 
its  freezing-point  is  one  of  far-reaching  consequences  in  the  opera- 
tions of  nature. 

When  a  mass  of  water,  such  as  a  lake,  is  exposed  to  the  influence 
of  a  cold  wind,  the  superficial  layer  of  water  is  cooled,  and  thereby 
becoming  specifically  denser,  it  sinks  to  the  bottom  and  exposes  a 
fresh  surface.  This  in  its  turn  has  its  temperature  lowered,  and  in 
like  manner  falls  to  the  bottom.  A  circulation  of  the  water  in  this 
way  is  set  up  until  the  entire  mass  reaches  a  temperature  of  4°. 
At  this  point  the  further  cooling  of  the  surface-layer  causes  expan- 
sion instead  of  contraction,  and  the  colder  water  becoming  speci- 
fically lighter  now  floats  upon  the  top,  where  it  remains  until  it 
congeals.  If  water  continued  to  contract  as  its  temperature  was 
reduced  below  4°,  the  circulatory  motion  would  continue  until  the 
whole  body  of  the  water  was  cooled  to  o°,  when  solidification  of  the 
entire  mass  would  take  place.  The  reason  that  certain  very  deep 

*  At  the  time  this  standard  was  first  adopted,  methods'  of  measurement  were 
less  refined  than  at  present.  In  reality  the  gramme  is  not  exactly  the  weight 
of  i  c.c.  of  water  at  its  point  of  maximum  density. 


2l6  Inorganic  Chemistry 

waters  seldom  or  never  freeze  is  because  the  duration  of  the  cold 
is  not  long  enough  to  bring  the  temperature  of  the  entire  mass 
of  the  water  down  to  4°,  and  until  that  is  effected  no  ice  can  form 
upon  the  surface. 

The  Solvent  Power  Of  Water.— Water  is  possessed  of  more 
general  solvent  powers  than  any  other  liquid  ;  that  is  to  say,  a  larger 
number  of  substances  are  dissolved  by  water  than  by  any  other 
liquid.  The  solvent  action  of  water  upon  gases,  liquids,  and  solids, 
in  so  far  as  it  is  shared  by  other  liquids,  has  been  dealt  with  under 
the  General  Properties  of  Liquids  (Part  I.,  chap.  xiii.). 

Water  Of  Crystallisation.— When  solid  substances  are  dis- 
solved in  water,  and  the  water  afterwards  evaporated,  the  dissolved 
substance  is  frequently  deposited  in  definite  crystalline  shapes. 
Many  salts  owe  their  crystalline  nature  to  the  fact  that  a  certain 
number  of  molecules  of  water  have  solidified  along  with  molecules 
of  the  salt,  each  molecule  of  the  salt  being  associated  with  a  defi- 
nite number  of  molecules  of  solid  water.  The  water  molecules 
must  be  regarded  as  having  entered  into  a  feeble  chemical  union 
with  the  salt  molecule,  but  a  union  which  is  of  a  somewhat  diffe- 
rent order  from  that  which  holds  together  the  atoms  of  oxygen  and 
hydrogen  in  the  water  molecules,  or  the  atoms  composing  the  salt 
in  the  salt  molecule  (see  p.  66).  Thus  copper  sulphate  crystallises 
associated  with  five  molecules  of  water,  CuSO4,5H2O  ;  magnesium 
sulphate  with  seven,  MgSO4,7H2O.  Water  so  associated  with, 
crystals  is  known  as  water  of  crystallisation,  and  the  compound 
is  called  a  hydrate. 

Many  salts  are  capable  of  crystallising  with  more  than  one  defi- 
nite number  of  molecules  of  water,  depending  upon  the  temperature 
at  which  the  crystallisation  takes  place  :  thus  sodium  carbonate, 
crystallised  at  the  ordinary  temperature,  has  the  composition 
Na2CO3,10H2O  ;  while  at  temperatures  between  30°  and  50°  the  salt 
that  is  deposited  contains  seven  molecules  of  water,  Na2CO3,7H2O. 
Sodium  chloride,  crystallised  from  solution  at  —  7°,  has  the  compo- 
sition, NaCl,2H2O  ;  while  the  crystals  that  are  deposited  at  —23°' 
contain  ten  molecules  of  water,  NaCl,10H2O. 

In  such  cases  as  these,  the  particular  crystalline  form  of  the  salt 
differs  with  the  different  degrees  of  hydration. 

Many  crystalline  salts,  when  exposed  to  the  air,  lose  either  some 
or  all  of  their  water  of  crystallisation,  and  in  so  doing  lose  their 
particular  geometric  form.  Thus  the  salt,  Na2CO3,10H2O  (ordinary 
washing  soda),  when  freely  exposed,  gradually  loses  its  crystalline 


Water  217 

form  and  falls  down  to  a  soft  white  powder,  which  consists  of  small 
crystals  of  another  form,  having  the  composition  Na2CO3,H2O. 
This  process  is  known  as  efflorescence,  the  crystals  being  said  to 
effloresce.  Other  crystals  undergo  exactly  the  reverse  change ;  they 
combine  with  moisture  from  the  air,  and  pass  into  other  crystalline 
forms  containing  more  water  of  crystallisation,  or  in  some  cases 
they  absorb  sufficient  moisture  to  cause  them  to  liquefy.  Such 
substances  are  said  to  deliquesce.  This  property  of  certain  salts  is 
made  use  of  for  withdrawing  traces  of  water  from  either  liquids  or 
gases.  Thus,  such  a  liquid  as  ether  may  be  freed  from  dissolved 
water  by  adding  to  it  copper  sulphate  containing  one  molecule  of 
water  of  crystallisation,  CuSO4,H2O ;  this  compound  takes  up  water 
and  passes  into  CuSO4,5H2O,  and  thereby  has  the  effect  of  drying 
the  ether.  Gases  in  the  same  way  are  frequently  dried  by  being 
passed  through  tubes  containing  calcium  chloride  from  which  the 
water  of  crystallisation  has  been  removed.  This  substance  absorbs 
water  with  avidity,  passing  into  the  hydrated  salt  CaCl2,6H2O. 

The  characteristic  colours  of  certain  salts  are  in  many  cases 
dependent  upon  the  amount  of  water  of  crystallisation  they  contain. 
Thus  cobalt  chloride,  CoCl2,6H2O,  is  a  pink  salt.  If  it  be  gently 
heated  to  120°  it  loses  its  water  and  becomes  CoCl2,  which  has  a 
rich  blue  colour.  Solutions  of  this  salt  have  been  employed  for 
the  so-called  sympathetic  inks.  The  faint  colour  of  the  pink  salt 
renders  words  written  upon  paper  with  its  dilute  solution  prac- 
tically invisible  ;  but  on  warming  the  paper,  and  thereby  expelling 
the  water  from  the  salt,  the  written  characters  appear  in  a  blue 
colour,  which  again  disappears  as  the  salt  is  allowed  to  rehydrate 
itself  by  exposure  to  the  air. 

One  of  the  most  striking  examples  of  this  change  of  colour 
resulting  from  varying  proportions  of  water  of  crystallisation  is 
seen  in  the  salt  magnesium  platino-cyanide,  which  crystallises  under 
ordinary  circumstances  as  a  bright  scarlet  salt  with  seven  molecules 
of  water,  MgPt(CN)4,7H2O.  When  this  salt  is  heated  to  about  50° 
it  loses  two  molecules  of  water,  and  is  converted  into  a  canary- 
yellow  salt,  MgPt(CN)4,5H2O.  If  the  temperature  be  raised  to 
1 00°  the  yellow  salt  becomes  white  by  the  loss  of  three  more  mole- 
cules, the  composition  of  the  white  salt  being  MgPt(CN)4,2H2O. 
When  a  solution  of  the  salt  is  carefully  evaporated  to  dryness  in 
a  dish  and  gently  warmed,  these  colour  changes  will  be  rendered 
evident  ;  and  upon  exposing  the  dried  and  white  residue  to  the  air, 
or  by  gently  breathing  into  the  dish,  the  salt  rehydrates  itself,  and 


21 8  Inorganic  Chemistry 

is  converted  into  the  crimson  compound  having  seven  molecules 
of  water. 

Many  salts  can  have  their  combined  water  withdrawn  by  power- 
ful dehydrating  agents  ;  thus,  if  a  crystal  of  copper  sulphate  ("blue 
vitriol/''  CuSO4,5H2O)  be  immersed  in  strong  sulphuric  acid,  the 
acid  abstracts  four  out  of  the  five  molecules  from  the  salt,  leaving 
the  nearly  white  salt  CuSO4,H2O  ;  or  when  alcohol  is  added  to  a 
solution  of  cobalt  chloride,  or  to  crystals  of  the  salt,  CoCl2,6H2O, 
the  alcohol  abstracts  water,  and  the  solution  becomes  blue. 

When  salts  containing  water  of  crystallisation  are  heated,  it 
frequently  happens  that  a  portion  of  the  water  is  more  easily  parted 
with  than  the  remainder.  Thus  copper  sulphate,  CuSO4,5H2O, 
when  heated  to  100°,  parts  with  four  molecules  of  water,  leaving  the 
salt  CuSO4,H2O  ;  and  in  order  to  drive  off  this  one  remaining  mole- 
cule, the  temperature  must  be  raised  above  200°.  Zinc  sulphate 
(or  white  vitriol),  ZnSO4,7H2O,  in  like  manner  loses  six  molecules  of 
water  at  100°,  but  retains  the  seventh  until  a  temperature  of  240°  is 
reached.  In  order,  therefore,  to  distinguish  between  the  water  that 
is  more  firmly  held  and  that  which  is  readily  parted  with,  the  term 
'water  Oj  ••  CD^titution  is  frequently  applied  to  the  former,  and  the 
fact  is  sometimes  expressed  in  notation  in  the  following  manner:— 

CuSO4H2O,4H2O ;  ZnSO4H2O,6H2O. 

Natural  Waters. — On  account  of  the  great  solvent  powers  of 
water,  this  compound  is  never  found  upon  the  earth  in  a  state  of 
absolute  purity  ;  even  rain,  as  it  falls  in  regions  far  removed  from 
the  dirty  atmosphere  of  towns,  not  only  dissolves  the  gases  of  the 
atmosphere,  but  also  small  quantities  of  those  suspended  matters 
which  are  always  present  in  the  air.  As  soon  as  the  rain  reaches 
the  earth,  the  water  at  once  exerts  its  solvent  action  upon  the 
mineral  matter  constituting  the  portion  of  the  earth's  crust  over 
which  it  flows,  and  through  which  it  percolates,  and  the  liquid  is 
rapidly  rendered  less  and  less  pure  as  it  travels  on  its  course  to 
lake  or  ocean. 

Natural  waters  may  be  broadly  divided  into  two  classes,  based 
upon  the  amount  of  dissolved  impurities  they  contain.  If  the  sub- 
stances in  solution  are  present  in  excessive  quantities,  or  to  such  an 
extent  as  to  be  perceptible  to  the  taste,  the  water  is  said  to  be  a 
mineral  water',  while,  on  the  other  hand,  waters  that  are  not  so 
rich  in  dissolved  impurities  are  known  as  fresh  waters. 


Natural  Waters  219 

Mineral  Waters. — The  most  exaggerated  examples  of  mineral 
waters  are  to  be  found  in  sea-water  and  in  the  waters  of  certain 
lakes,  which,  having  no  outlet,  are  fulfilling  the  purpose  of  enormous 
evaporating  basins,  in  which  the  waters  that  flow  into  them  are 
undergoing  evaporation  and  therefore  concentration  ;  such,  for 
example,  as  the  salt  lakes  of  Egypt,  the  Elton  lake  in  Russia,  and 
the  Dead  Sea.  In  waters  of  this  description  the  total  quantity  of 
dissolved  solid  matter  is  very  considerable,  and,  as  in  the  case  of 
the  Dead  Sea,  is  often  deposited  in  crystalline  masses  round  the 
shores  of  the  lake.  The  following  table  gives  the  total  amount  of 
dissolved  saline  matter  contained  in  1000  grammes  of  certain  of 
these  waters  : — 

Irish  Sea 33.86 

Mediterranean  Sea         .         .        .       40.0 

Dead  Sea 228.57 

Elton  Lake 271.43 

As  a  typical  example  of  a  sea  water,  the  composition  of  the 
water  of  the  English  Channel  may  be  quoted  ;  1000  grammes  of 
this  water  contain — 

Sodium  chloride 27.059 

Magnesium  chloride        ....  3.666 
Magnesium  sulphate        .        .        .        .       2.296 

Calcium  sulphate 1.406 

Potassium  chloride  .....  0.766 

Calcium  carbonate 0.033 

Magnesium  bromide        .  0.029 

35.255 

Water 964-745 

1000.000 

Passing  from  these  highly  concentrated  mineral  waters,  we  find 
a  large  number  of  spring  waters  which  are  classed  as  mineral,  not 
because  the  total  quantity  of  foreign  matter  in  solution  is  excessive, 
but  rather  because  they  contain  an  abnormally  large  proportion  o{ 
a  few  special  substances.  Thus,  large  quantities  of  magnesium 
sulphate  and  chloride  are  found  in  such  springs  as  those  at 
Epsom  and  Friedrichshall.  Others  are  found  to  contain  consider- 
able quantities  of  sodium  sulphate  and  sodium  carbonate  ;  while 


22O  Inorganic  Chemistry 

those  known  as  chalybeate  waters  contain  ferrous  carbonate  in 
solution.  Spring  waters  that  are  charged  with  unusual  quantities  of 
soluble  gases  are  likewise  placed  in  the  category  of  mineral  waters, 
such  as  the  waters  of  Apollinaris  and  Seltzer,  containing  large 
quantities  of  carbon  dioxide  ;  and  the  sulphur  springs  at  Harrogate 
and  Aachen,  which  hold  in  solution  sulphuretted  hydrogen  as  well 
as  alkaline  sulphides. 

Fresh  Waters. — The  purest  form  of  natural  water  is  rain-water. 
The  average  weight  of  solid  matter  dissolved  in  rain-water,  col- 
lected in  the  country  and  in  perfectly  clean  vessels  upon  which  it 
exerts  no  solvent  action,  is  found  to  be  0.0295  Part  m  IOO°  parts 
of  water.  Collected  in  or  near  towns,  rain-water  always  contains 
a  larger  amount  of  dissolved  impurities,  such  as  nitrates,  sulphates, 
ammoniacal  salts,  and  often  considerable  quantities  of  sulphuric 
acid  :  it  is  the  acid  nature  of  the  rain  that  causes  so  much  damage 
to  stone  buildings. 

The  nature  and  extent  of  the  contamination  that  rain-water 
suffers  after  it  has  fallen  must  obviously  depend  very  largely  upon 
geographical  and  geological  circumstances,  and  therefore  there  are 
no  special  features  that  are  distinctly  characteristic  of  waters  from 
rivers,  lakes,  or  springs. 

Thus,  the  total  solid  impurity  in  1000  parts  of  water  from  the 
river  Dee  at  Aberdeen  is  0.057,  while  that  contained  in  the 
Thames  is  0.30  parts. 

The  water  of  Loch  Katrine  only  contains  0.032  part  of  solid 
matter  dissolved  in  1000  parts,  while  that  of  Elton  lake  contains 
as  much  as  271.43. 

The  same  wide  differences  are  also  seen  in  spring  waters  from 
different  geological  strata.  Spring  waters  from  granite  and  gneiss 
rocks  contain  on  an  average  0.059  part  of  dissolved  solid  matter 
in  looo  parts,  while  those  from  magnesian  limestone  average  as 
much  as  0.665  part.  As  a  broad  general  rule,  river  waters  are 
found  to  contain  less  solid  matter  in  solution  than  spring  waters, 
and  these  in  their  turn  less  than  deep  well  waters.  Thus,  com- 
paring waters  from  different  sources,  and  selecting  only  such 
samples  as  are  known  to  be  free  from  pollution  from  either  sewage 
matter  or  other  abnormal  impurities,  it  will  be  seen  that,  with 
regard  to  the  dissolved  solid  matter  they  contain,  they  fall  in  the 
following  order  : — 


Natural  Waters  221 

Total  Solid  Impurity  Dissolved  in  1000  Parts  of 
Unpolluted  Waters. 

Rain-water  (average  of  39  samples)          .  .  .0295 

Rivers  and  lakes  (average  of  195  samples)  .  .0967 

Spring  waters  (average  of  198  samples)  .  .  .2820 

Deep  well  waters  (average  of  157  samples)  .  .4378 

Hardness  Of  Water.— Certain  of  the  salts  that  are  very  fre- 
quently present  as  impurities  in  natural  waters  give  to  these 
waters  the  property  that  is  known  as  hardness.  The  chief  com- 
pounds that  produce  this  effect  are  the  salts  of  calcium  and 
magnesium.  The  term  hardness  is  applied  to  such  waters  on 
account  of  the  difficulty  of  obtaining  a  lather,  with  soap,  in  the 
ordinary  process  of  washing.  Pure  soap  may  be  regarded  as  a 
mixture  of  the  sodium  salts  of  certain  fatty  acids  (oleic,  stearic, 
palmitic,  £c.),  which  are  soluble  in  pure  water.  In  the  presence 
of  salts  of  calcium  and  magnesium  the  soap  is  decomposed,  and 
an  insoluble  curdy  precipitate  is  formed  by  the  union  of  the  fatty 
acid  of  the  soap  with  the  calcium  and  magnesium  of  the  salts. 
Until  the  whole  of  the  hardening  salts  have  in  this  way  been 
thrown  out  of  solution,  no  lather  can  be  obtained,  and  the  soap  is 
useless  as  a  cleansing  agent ;  but  as  soon  as  this  point  is  reached, 
the  addition  of  any  further  quantity  of  soap  at  once  raises  a  lather 
on  the  water,  and  the  soap  is  capable  of  acting  as  a  detergent. 
This  process  of  precipitating  the  salts  of  calcium  and  magnesium 
is  known  as  softening,  and  in  this  instance  the  water  is  softened  at 
the  expense  of  the  soap. 

Hard  waters  often  become  less  hard  after  being  boiled  for  a 
short  time,  and  this  hardness  which  is  so  removed  is  termed  the 
temporary  hardness.  The  degree  of  hardness  which  the  water  still 
possesses  after  prolonged  boiling  is  distinguished  by  the  term 
permanent  hardness.  The  diminution  of  the  total  hardness  of 'a 
water  by  boiling  is  due  to  the  fact  that  the  soluble  acid  carbonates 
of  calcium  and  magnesium  are  decomposed  during  this  process 
into  water,  carbon  dioxide  (which  escapes  as  gas),  and  the  prac- 
tically insoluble  normal  carbonates  of  these  metals  ;  thus,  in  the 
case  of  the  calcium  salt — 

CaH2(CO3)2=H2O  +  CO2  +  CaCO3. 

When  such  a  water  is  boiled,  the  calcium  carbonate  is  thrown 
down  as  a  white  precipitate,  which  gradually  collects  upon  the 


222  Inorganic  Chemistry 

bottom  of  the  containing  vessel.  The  "  furring  "  of  kettles,  and  the 
formation  of  calcareous  deposits  in  boilers,  is  largely  due  to  this 
cause. 

In  the  case  of  waters  that  are  highly  charged  with  calcium  car- 
bonate, held  in  solution  by  dissolved  carbonic  acid,  this  deposition 
of  calcium  carbonate  may  even  take  place  at  the  ordinary  tempe- 
rature, owing  to  the  diffusion  of  the  dissolved  carbon  dioxide  into 
the  air.  It  is  in  this  way  that  those  remarkable,  and  often  beauti- 
fully fantastic  formations,  known  as  stalactites,  have  been  produced 
in  certain  subterranean  caves.  Water  charged  with  the  soluble 
calcium  carbonate,  in  slowly  dropping  from  the  roof  of  such  a  cave, 
loses  a  portion  of  its  dissolved  carbon  dioxide,  and,  in  consequence, 
deposits  a  certain  amount  of  the  calcium  carbonate  which  was  in 
solution.  Each  drop,  as  it  slowly  forms,  adds  its  little  share  of 
calcium  carbonate  to  the  deposit,  which  thereby  gradually  grows, 
much  as  an  icicle  grows,  as  a  dependent  mass  called  a  stalactite. 
Whether  the  water  that  drops  from  the  stalactite  has  deposited 
the  whole  of  its  calcium  carbonate,  will  depend  largely  upon  the 
time  occupied  by  each  drop  in  gathering  and  dropping  ;  if,  as  often 
happens,  the  whole  has  not  been  precipitated,  the  remainder  is 
deposited  upon  the  floor  of  the  cave,  and  a  growing  column  of 
calcium  carbonate,  called  a  stalagmite,  gradually  rises  from  the 
ground  until  it  ultimately  meets  the  stalactite. 

Clark's  Process  for  Softening-  Water.— Waters  whose  hard- 
ness is  due  to  the  presence  of  the  carbonates  of  calcium  and 
magnesium  can  be  deprived  of  their  hardness  by  the  addition  to 
them  of  lime.  The  amount  of  hardness  is  first  estimated,  and  such 
an  amount  of  milk  of  lime  is  then  added  as  is  demanded  by  the 
following  equation  : — 

CaH2(CO3)2  +  CaO  =  H2O  +  2CaCO3. 

In  this  way  the  soluble  calcium  salt  is  converted  into  the  insoluble 
normal  carbonate,  which  settles  to  the  bottom  of  the  tank. 

The  salts,  which  are  mainly  instrumental  in  causing  the  per- 
manent hardness,  are  the  sulphates  of  calcium  and  magnesium. 
The  degree  of  hardness  and  its  particular  order,  that  is,  whether 
temporary  or  permanent,  will  obviously  be  determined  entirely  by 
the  particular  geological  formation  from  which  the  water  is  derived. 

Potable  Waters.— Undoubtedly  the  most  important  use  to 
which  water  is  put  is  its  employment  as  an  article  of  food  to  man, 


Natural  Waters  223 

and  since  it  has  been  proved  beyond  dispute  that  many  virulent 
diseases,  such  as  cholera,  typhoid  fever,  and  others,  are  propagated 
through  the  medium  of  drinking-water,  it  becomes  a  matter  of  the 
greatest  sanitary  importance  that  the  waters  supplied  for  this  pur- 
pose should  be  as  pure  as  possible.  Excepting  in  very  rare  in- 
stances, where  poisonous  mineral  matters  accidentally  gain  access 
to  drinking-water  (as,  for  example,  in  the  case  of  certain  waters 
which  are  capable  of  attacking,  and  to  a  slight  extent  dissolving, 
the  lead  of  the  pipes  through  which  they  may  be  passed),  the  solid 
matters  that  are  usually  found  in  waters  are  not  injurious  to  health. 
The  living  germs  or  bacilli,  through  whose  agency  zymotic  diseases 
are  caused,  cannot  be  detected  in  a  sample  of  water  by  any  direct 
chemical  analysis.  A  specimen  of  pure  distilled  water  might 
be  artificially  contaminated  with  such  organisms  so  as  to  con- 
stitute it  a  most  virulent  poison,  and  still  chemical  analysis 
would  fail  to  detect  the  danger,  and  the  water  would  be  pronounced 
pure.  Chemical  analysis  can,  however,  reveal  the  presence  of 
excrementitious  matter,  and  also  of  the  characteristic  products  re- 
sulting from  its  decomposition :  it  can  with  certainty  detect  in  the 
water  the  evidence  of  recent  contamination  with  sewage  matters, 
and  it  can  also,  with  considerable  precision,  trace  the  evidences 
of  its  having  been  so  contaminated  at  an  earlier  stage  of  its  history. 
It  cannot,  however,  distinguish  between  pollution  with  healthy  and 
with  infected  excreta,  and  therefore  it  is  necessary  to  regard  with 
the  greatest  suspicion  any  water  to  which  sewage  has  at  any  time 
gained  access.  Waters  that  are  made  use  of  for  drinking  purposes 
may  be  classified  in  the  following  order  : — 

(  I.  Spring  water. 

Safe      .     .  <  2.  Deep  well  water. 

(  3.  Mountain  rivers  and  lakes. 

q       .  .  (  4.  Stored  rain-water. 

(  5.  Surface  water  from  cultivated  land. 

(  6.  River  water  to  which  sewage  gains  access. 
Dangerous    <  „ 

7.  Shallow  well  water. 


HYDROGEN  PEROXIDE. 

Formula, 


Occurrence.  —  This  compound  is  occasionally  found  in  small 
quantities  in  the  atmosphere,  and  also  in  dew  and  rain. 


224  Inorganic  Chemistry 

Modes  Of  Formation.—  (  i.)  Hydrogen  peroxide  is  produced  in 
small  quantities  during  the  burning  of  hydrogen  in  the  air.  If  a 
jet  of  burning  hydrogen  be  caused  to  impinge  upon  the  surface  of 
water,  the  temperature  of  which  is  not  allowed  to  rise  above  20°, 
the  water  will  be  found,  after  a  short  time,  to  contain  hydrogen 
peroxide.* 

(2.)  This  compound  is  also  produced  by  the  decomposition  of 
barium  peroxide  by  carbonic  acid.  For  this  purpose  a  stream  of 
carbon  dioxide  is  passed  through  ice-cold  water,  into  which  from 
time  to  time  small  quantities  of  barium  peroxide  are  stirred. 
Barium  carbonate  is  precipitated,  and  a  dilute  aqueous  solution  of 
hydrogen  peroxide  is  obtained  — 

BaO2  +  H2CO3  =  BaCO3  +  H2O2. 

(3.)  Barium  peroxide  may  be  decomposed  by  either  hydrochloric, 
sulphuric,  silicofluoric,  or  phosphoric  acid.  Whichever  acid  be 
employed,  the  barium  peroxide,  previously  mixed  with  a  small 
quantity  of  water,  is  added  gradually  to  the  acid  ;  which,  in  the  case 
of  either  hydrochloric  or  sulphuric  acid,  should  be  diluted  with  from 
five  to  ten  times  its  volume  of  water.  The  temperature  of  the 
mixture  is  not  allowed  to  rise  above  20°.  Thus,  in  the  case  of 
hydrochloric  acid  — 

=  BaCl2+H202, 


the  soluble  barium  chloride  is  removed  by  the  addition  of  sulphuric 
acid,  whereby  barium  sulphate  is  precipitated  and  hydrochloric 
acid  formed  — 


The  hydrochloric  acid  may  be  removed  by  adding  a  solution  of 
silver  sulphate,  which  precipitates  silver  chloride,  leaving  sulphuric 
acid  in  solution  — 

2HC1  +  Ag2SO4=2AgCI  +  H2SO4. 

And,  lastly,  the  free  sulphuric  acid  is  withdrawn  by  the  addition  of 
barium  carbonate  — 

H2SO4  +  BaCO3=  BaSO4  +  H2O  +  CO2. 

When  sulphuric  acid  is  employed  for  the  decomposition  of  barium 
peroxide,  the  crystallised,  or  hydrated  peroxide  (BaO2,8H2O),  is 

*  See  "  Chemical  Lecture  Experiments,"  new  ed.,  p.  74. 


Hydrogen  Peroxide  225 

most  advantageous  for  the  purpose.  This  salt,  made  into  a  paste 
with  water,  is  gradually  added  to  the  diluted  and  cooled  acid,  until 
the  acid  is  nearly  but  not  quite  neutralised.  The  slight  excess  of 
acid  is  removed  by  the  addition  of  the  exact  quantity  of  barium 
hydroxide  (baryta-water)  necessary  to  neutralise  it,  and  the  insoluble 
barium  sulphate  is  removed  by  filtration.  On  a  large  scale  silico- 
fluoric  acid  or  phosphoric  acid  is  usually  employed,  preferably  the 
latter,  as  it  is  found  that  small  quantities  of  free  phosphoric  acid 
in  hydrogen  peroxide  greatly  retard  its  decomposition. 

(4.)  Hydrogen  peroxide  is  also  readily  obtained  by  decomposing 
potassium  peroxide  by  means  of  tartaric  acid.  The  potassium 
peroxide  is  added  to  a  cooled  strong  aqueous  solution  of  tartaric 
acid,  when  potassium  tartrate  separates  out,  and  an  aqueous  solu- 
tion of  hydrogen  peroxide  is  obtained. 

(5.)  When  small  quantities  of  hydrogen  peroxide  are  required 
for  the  purpose  of  illustrating  its  properties,  it  is  most  conveniently 
obtained  by  adding  sodium  peroxide  to  dilute  and  well-cooled 
hydrochloric  acid,  whereby  sodium  chloride  and  hydrogen  per- 
oxide are  formed,  both  of  which  remain  in  solution  — 


Na2O2  +  2HCl  =  2NaCl  +  H2O2. 

(6.)  Hydrogen  peroxide  is  formed  in  considerable  quantity  when 
ozone  is  passed  through  ether  floating  upon  water.  Probably  a 
peroxidised  compound  of  ether  is  first  produced,  which  is  then 
decomposed  by  the  water.  This  production  of  hydrogen  peroxide 
may  readily  be  demonstrated  by  placing  a  small  quantity  of  water 
and  ether  in  a  beaker,  and  suspending  into  the  vapour  a  spiral  of 
platinum  wire  which  has  been  gently  heated.  The  combustion  of 
the  ether  vapour  upon  the  wire,  whereby  the  latter  is  maintained 
at  a  red  heat,  is  attended  with  the  formation  of  ozone,  and  this 
acting  upon  the  ether,  as  already  described,  results  in  the  pro- 
duction of  hydrogen  peroxide,  which  may  be  detected  in  solution 
in  the  water. 

(7.)  In  small  quantities,  hydrogen  peroxide  is  produced  when 
moist  ether  is  exposed  to  the  action  of  oxygen,  under  the  prolonged 
influence  of  sunlight. 

Properties.—  The  dilute  aqueous  solution  of  hydrogen  peroxide, 
obtained  by  the  foregoing  methods,  is  concentrated  by  evaporation 
over  sulphuric  acid  in  vacuo.  In  the  pure  condition  it  is  a  colour- 
less and  odourless,  syrupy  liquid,  having  an  extremely  bitter  and 

P 


226  Inorganic  Chemistry 

metallic  taste.  The  specific  gravity  of  the  liquid  is  1.4532.  The 
substance  is  extremely  unstable,  giving  up  some  of  its  oxygen  even 
at  temperatures  as  low  as  -  20°,  and  decomposing  with  explosive 
violence  when  heated  to  100°.  Hydrogen  peroxide  bleaches 
organic  colours,  but  less  rapidly  than  chlorine.  When  placed 
upon  the  skin  it  destroys  the  colour,  and  gives  rise  to  an  irritating 
blister.  When  diluted  with  water,  and  especially  if  rendered  acid, 
the  compound  is  far  more  stable,  and  in  this  condition  may  be 
preserved  at  the  ordinary  temperature  for  a  considerable  length  of 
time.  When  such  an  aqueous  solution  is  strongly  cooled,  it  deposits 
ice,  and  in  this  way,  by  the  removal  of  the  frozen  water,  the  solu- 
tion may  be  concentrated.  Hydrogen  peroxide  itself  solidifies 
between  —20°  and  —23°.  When  heated  the  solution  is  decom- 
posed into  water  and  oxygen — 

H2O2=H2O  +  O. 

Owing  to  the  readiness  with  which  hydrogen  peroxide  gives  up 
the  half  of  its  oxygen  and  is  converted  into  water,  its  properties 
are  generally  those  of  a  powerful  oxidising  agent.  It  liberates 
iodine  from  potassium  iodide  ;  it  converts  sulphurous  acid  into 
sulphuric  acid,  and  oxidises  lead  sulphide  into  lead  sulphate.  Its 
action  upon  lead  sulphide  is  made  use  of  in  restoring  something 
of  the  original  brilliancy  to  oil  paintings  that  have  become  dis- 
coloured. The  "  white-lead "  used  in  oil  paints  is  gradually  con- 
verted into  lead  sulphide  when  such  paintings  are  exposed  to  air, 
especially  the  air  of  towns,  which  is  liable  to  contain  small 
quantities  of  sulphuretted  hydrogen.  Lead  sulphide  being  black, 
the  picture  slowly  assumes  a  uniformly  dark  colour.  When  such 
a  discoloured  picture  is  washed  over  with  dilute  hydrogen  peroxide, 
the  black  sulphide  is  oxidised  into  the  white  lead  sulphate — 

PbS  +  4H2O2  =  4H2O  +  PbSO4. 

This  compound  is  employed  for  bleaching  articles  that  would 
suffer  injury  by  the  use  of  other  bleaching  agents,  such  as  ivory, 
feathers,  and  even  the  teeth. 

Hydrogen  peroxide  is  also  capable  of  oxidising  hydrogen. 
Thus  when  a  dilute  acidulated  solution  of  the  peroxide  is  electro- 
lysed, oxygen  is  evolved  at  the  anode,  but  no  gas  escapes  from 
the  cathode  ;  the  nascent  hydrogen  being  oxidised  to  water — 


Hydrogen  Peroxide  227 

Hydrogen  peroxide,  in  many  of  its  reactions,  appears  to  act  as  a 
deoxidising  agent  ;  thus,  manganese  dioxide  in  contact  with  this 
substance  is  reduced  to  mahganous  oxide  — 

MnO2  +  H2O2=  MnO  +  O2  4-  H2O. 

Similarly,  silver  oxide  is  reduced  to  metallic  silver  with  the 
evolution  of  oxygen  — 


In  like  manner,  when  ozone  is  acted  upon  by  hydrogen  per- 
oxide, a  reaction  takes  place  exactly  analogous  to  that  with  silver 
oxide,  which  will  be  the  more  obvious  if  the  formula  for  ozone  be 
written  O2O  instead  of  O3,  thus  — 


Although,  in  a  sense,  these  reactions  may  be  regarded  as  reduc- 
ing,  or  deoxidising,  actions,  in  essence  they  are  not  different  from 
those  which  have  been  given  as  illustrative  of  the  oxidising  power 
of  hydrogen  peroxide.  It  will  be  seen  that  they  all  depend  upon 
the  readiness  with  which  the  compound  parts  with  an  atom  of 
oxygen,  but  that  in  these  latter  cases  the  oxygen  that  is  so  given 
up  is  engaged  in  oxidising  another  atom  of  oxygen,  contained  in  the 
other  compound.  Thus,  in  the  case  of  silver  oxide,  its  atom  of 
oxygen  is  oxidised  by  the  liberated  oxygen  from  the  hydrogen 
peroxide,  and  converted  into  the  complete  molecule  of  oxygen. 
By  these  reactions  Brodie  first  demonstrated  the  dual,  or  di- 
atomic, character  of  the  molecule  of  oxygen. 

When  hydrogen  peroxide  is  added  to  a  dilute  acidulated  solution 
of  potassium  dichromate,  a  deep  azure-blue  solution  is  obtained 
(see  Chromium),  which  affords  a  delicate  test  for  this  com- 
pound. To  apply  the  test,  the  dilute  hydrogen  peroxide  is  shaken 
up  with  ether,  a  few  drops  of  acidulated  potassium  dichromate 
are  then  added,  and  the  mixture  again  shaken.  The  blue  com- 
pound being  more  soluble  in  ether  than  in  water,  the  ethereal 
liquid  will  separate  as  a  blue  layer.  In  this  way,  the  presence  of 
0.00025  grammes  of  hydrogen  peroxide  in  20  c.c.  of  water  can  be 
detected. 

Hydrogen  peroxide  is  decomposed  by  contact  with  many  sub- 
stances which  themselves  do  not  combine  with  the  oxygen  ;  thus 
charcoal,  finely  divided  palladium,  platinum,  mercury,  and  notably 
silver,  when  brought  into  hydrogen  peroxide,  determine  its  decom- 


228  Inorganic  Chemistry 

position  into  water  and  oxygen,  the  rapidity  of  the  action  being 
increased  if  the  liquid  be  made  alkaline.  The  action  is  doubtless 
catalytic,  although  in  all  cases  the  exact  modus  operandi  is  not 
clearly  understood.  In  the  case  of  silver  it  is  believed  that  silver 
oxide  (perhaps  peroxide)  is  first  formed,  and  then  decomposed, 


Ag2O  +  H2O2  =  H2O  +  O2  +  Ag2. 

When  hydrogen  peroxide  is  added  to  solutions  of  the  hydroxides 
of  barium,  strontium,  or  calcium,  the  peroxide  of  the  metal  is 
precipitated  — 


The  compound  is  deposited  in  crystals  having  the  composition 
BaO2,8H2O. 

With  the  hydroxides  of  the  alkali  metals,  the  peroxide  (which  is 
soluble  in  water)  may  be  precipitated  by  the  addition  of  alcohol  ; 
when  in  the  case  of  sodium  peroxide,  crystals  are  obtained  of 
Na.2O2,8H2O. 

Hydrogen  peroxide  is  a  useful  antiseptic  ;  it  possesses  the  ad- 
vantages of  being  free  from  smell,  without  poisonous  or  injurious 
action  upon  the  system,  and  of  leaving  as  a  residue,  after  having 
furnished  its  available  oxygen,  only  water. 

The  constitution  of  hydrogen  peroxide  is  usually  expressed  by  the  formula 
H-O-O-H,  but  the  accumulating  evidence  that  oxygen  is  capable  of  functioning 
as  a  quadrivalent  element  has  led  to  the  view  that  its  constitution  is  better 

H\ 
represented  by  the  formula        /O  :  O. 

Those  metallic  peroxides  which  yield  hydrogen  peroxide  on  treatment  with 
dilute  acids  may  also  be  regarded  as  similarly  constituted,  thus 
Nax. 

)>O  :  O  and  Ba  :  O  :  O 
Na/ 

Ba:O:O  +  2HCl  =  BaCl2  +  H2O:O 

while  those  peroxides  which  yield  oxygen  (or  its  equivalent  of  chlorine)  under 
similar  treatment,  contain  only  divalent  oxygen  atoms,  e.g. 

O  :  Pb  :  O  +  2HC1  =  PbCl2  +  H2O  +  O. 
Ozone  may  be  regarded  as  peroxide  of  oxygen,  and  expressed  by  the  formula 

O  :  O  :  O  instead  of  the  more  familiar  formula  \    /  in  which  each  atom  is 

O 

represented  as  being  divalent. 


CHAPTER    IV 
NITROGEN 

Symbol,  N.     Atomic  weight  «=  14.01.     Molecular  weight  =  28.02. 

History. — Nitrogen  was  discovered  by  Rutherford  in  1772.  He 
showed  that  when  an  animal  is  placed  in  a  confined  volume  of  air 
for  some  time,  and  the  air  afterwards  treated  with  caustic  potash, 
to  absorb  from  it  the  carbon  dioxide  ("fixed  air"),  there  still 
remained  a  gas  which  was  incapable  of  supporting  either  respira- 
tion or  combustion.  He  called  the  gas  mephitic  air.  Scheele  was 
the  first  to  recognise  that  this  gas  was  a  constituent  of  the  air. 
Lavoisier  applied  the  name  azote  to  the  gas,  to  denote  its  inability 
to  support  life.  The  name  nitrogen,  signifying  the  nitre-producer, 
was  suggested  by  Chaptal,  from  the  fact  that  the  gas  was  a  con- 
stituent of  nitre. 

Occurrence. — In  the  free  state  nitrogen  is  present  in  the  atmos- 
phere, of  which  it  forms  about  four-fifths.  Certain  nebulas  have 
been  shown  by  spectroscopic  observation  to  contain  nitrogen  in 
the  uncombined  condition.  In  combination,  nitrogen  is  found  in 
ammonia,  in  nitre  (potassium  nitrate),  and  in  a  great  number  of 
animal  and  vegetable  compounds. 

Modes  of  Formation.— ( i.)  Nitrogen  is  very  readily  obtained 
from  the  atmosphere  by  the  abstraction  of  the  oxygen  with  which 
it  is  there  mixed.*  This  is  conveniently  done  by  burning  a  piece 
of  phosphorus  in  air,  confined  over  water.  The  phosphorus  in 
burning  combines  with  the  oxygen,  forming  dense  white  fumes  of 
phosphorus  pentoxide,  which  gradually  dissolve  in  the  water,  and 
nitrogen  remains  in  the  vessel.  The  nitrogen  obtained  in  this  way 
is  never  quite  pure,  for  the  phosphorus  becomes  extinguished 
before  the  oxygen  is  entirely  removed.  It  is  also  admixed  with 
the  other  gases  present  in  the  atmosphere  (argon,  carbon  dioxide, 
&c.  ;  see  Atmosphere),  amounting  in  all  to  about  I  per  cent,  of 
the  total. 

(2.)  Nitrogen  in  a  purer  state  can  be  prepared  from  the  atmos- 

*  Experiments  254,  255,  "  Chemical  Lecture  Experiments,"  new  ed. 

229 


230  Inorganic  Chemistry 

phere  by  passing  a  stream  of  pure  air  over  metallic  copper  con- 
tained in  a  combustion  tube,  and  heated  to  redness  in  a  furnace. 
The  air  is  contained  in  a  gas-holder,  and  is  passed  through  two 
U  -tubes,  the  first  containing  potassium  hydroxide  (caustic  potash), 
in  order  to  absorb  the  carbon  dioxide  ;  and  the  second  filled  with 
fragments  of  pumice  moistened  with  sulphuric  acid,  in  order  to 
arrest  the  aqueous  vapour.  The  purified  air,  on  passing  over  the 
heated  copper,  is  deprived  of  the  whole  of  its  oxygen,  cupric  oxide, 
CuO,  being  formed,  while  the  nitrogen  passes  on  and  may  be 
collected.  This  gas  contains  small  quantities  of  argon  (p.  256). 

(3.)  Oxygen  is  rapidly  absorbed  by  a  solution  of  cuprous  chloride 
in  hydrochloric  acid  ;  a  ready  method,  therefore,  of  obtaining 
nitrogen  from  the  air  is  to  place  a  quantity  of  this  solution  in  a 
stoppered  bottle,  and  shake  it  up  with  the  contained  air.  The 
colourless  cuprous  chloride  solution  quickly  absorbs  the  oxygen, 
becoming  dark  in  colour,  and  being  converted  into  cupric  chloride, 
the  nitrogen  of  the  air  remaining  in  the  bottle  — 


(4.)  Nitrogen  is  obtained  by  heating  a  strong  solution  of  ammo- 
nium nitrite  in  a  flask,  the  salt  splitting  up  into  water  and  nitrogen  — 

NH4NO2  =  2H2O  +  N2. 

In  practice  it  is  found  more  convenient  to  employ  a  mixture  of 
ammonium  chloride  and  sodium  nitrite  — 

NH4Cl  +  NaNO2=NaCl  +  2H2O  +  N2. 

(5.)  By  heating  a  mixture  of  ammonium  nitrate  and  ammonium 
chloride,  a  mixture  of  nitrogen  and  chlorine  is  evolved  ;  the  latter 
gas  may  be  absorbed,  by  passing  the  mixture  through  either  milk 
of  lime  or  a  solution  of  sodium  hydroxide  — 

2NH4NO3+NH4C1  =  5N  +  C1  +  6H2O. 

(6.)  Nitrogen  is  also  evolved  when  ammonium  chromate,  or  a 
mixture  of  potassium  dichromate  and  ammonium  chloride,  is 
heated— 

(N  H4)2Cr2Or  =  Cr2O3  +  4H2O  +  N2, 
or— 


Nitrogen 


231 


(7.)  When  ammonia  is  acted  upon  by  chlorine  it  is  decomposed, 
the  chlorine  combining  with  the  hydrogen  to  form  hydrochloric 
acid,  and  the  nitrogen  being  liberated  — 


If  the  chlorine  be  passed  into  a  strong  solution  of  ammonia,  the 
hydrochloric  acid  which  is  produced  combines  with  the  excess  of 
ammonia,  forming  ammonium  chloride  ;  thus  — 


The  chlorine,  after  being  washed  by  passing  through  water,  is 
bubbled  through  strong  aqueous  ammonia  contained  in  a  Woulf  s 
bottle.  As  each  bubble  of  chlorine  enters  into  the  ammonia,  the 


FIG.  47. 

combination  is  attended  by  a  feeble  yellowish  flash  of  light,  and  a 
rapid  stream  of  nitrogen  is  evolved.  The  nitrogen,  which  carries 
with  it  dense  white  fumes  of  ammonium  chloride,  should  be  scrubbed 
by  being  passed  through  a  second  bottle  filled  with  fragments  of 
broken  glass  moistened  with  water,  and  it  can  then  be  collected 
over  water  in  the  ordinary  way,  as  shown  in  Fig.  47.*  In  prepar- 
ing nitrogen  by  this  reaction  it  is  very  necessary  that  the  ammonia 
should  be-in  considerable  excess,  otherwise  there  is  liable  to  be 
formed  the  dangerously  explosive  compound  of  nitrogen  and  chlo- 
rine (see  Nitrogen  Trichloride). 

*  Experiment  261. 


232 


Inorganic  Chemistry 


Properties. — Nitrogen  is  a  colourless  gas  without  taste  or 
smell.  It  is  slightly  Hghter  than  air,  its  specific  gravity  being 
0.973  (air  =  i).  One  litre  of  the  gas  at  o°  C.  and  760  mm.  weighs 
14  criths,  or  1.250  grammes. 

Nitrogen  is  only  very  slightly  soluble  in  water,  its  coefficient  of 
absorption  at  o°  C.  being  0.020346. 

Nitrogen  will  not  burn,  neither  will  it  support  the  combustion  of 
ordinary  combustibles.  It  is  not 
poisonous,  but  is  incapable  of  sup- 
porting respiration. 

Nitrogen  is  an  extremely  inert 
substance,  combining  directly,  and 
with  difficulty,  with  only  a  very  few 
elements.  Under  the  influence  of 
the  high  temperature  of  the  electric 
spark  it  can  be  made  to  unite  directly 
with  oxygen  (see  p.  235).  Certain 
metals  also  combine  directly  with  it, 
forming  nitrides.  Thus,  when  lithium 
or  magnesium  are  heated  in  nitrogen, 
they  form  respectively  NLi3  and 
N2Mg3.  This  reaction  may  be  con- 
veniently shown  by  means  of  the 
apparatus  seen  in  Fig.  48.  A  small 
quantity  of  powdered  magnesium  is 
placed  in  a  hard  glass  tube,  which 
is  connected  to  a  long  narrow  tube 
dipping  into  water,  and  a  stream  of 
nitrogen  is  passed  through.  When 
the  air  is  all  displaced  the  passage  of 

the  nitrogen  is  stopped  and  the  magnesium  strongly  heated.  At  a 
red  heat  the  nitrogen  will  be  rapidly  absorbed,  and  the  water  will 
be  seen  to  rise  in  the  long  tube. 

This  property  of  nitrogen  of  uniting  directly  with  magnesium 
was  utilised  in  effecting  the  separation  of  the  nitrogen  of  the  air 
from  the  small  quantities  of  argon  and  other  "  inert  gases  "  con- 
tained in  the  atmosphere. 

Although  it  is  true  that  nitrogen  in  the  elemental  condition  is  an 
inert  substance,  the  element  itself  is  in  reality  possessed  of  strong 
chemical  affinities.  Indeed,  the  very  inertness  of  its  molecules 
may  be  regarded  as  an  indication  of  the  strong  affinity  between 


FIG.  48. 


Nitrogen  233 

the  two  atoms  which  constitute  the  molecule.  Nitrogen  enters 
into  the  composition  of  an  enormous  number  of  compounds,  and 
its  atoms  must  be  regarded  as  possessing  great  chemical  activity. 
The  formation  of  such  compounds  as  the  nitrides  above  mentioned 
may  be  quoted  as  an  illustration.  Although  elementary  nitrogen 
combines  directly  with  comparatively  few  metals,  and  with  most 
of  these  only  at  somewhat  high  temperatures,  these  compounds 
are  readily  produced  if,  instead  of  elementary  nitrogen,  nitro- 
gen in  combination  with  hydrogen  (ammonia)  be  employed  (see 
P-  278). 

The  critical  temperature  of  nitrogen  is  -  149°,  and  when  cooled 
to  this  point  a  pressure  of  27.5  atmospheres  causes  its  liquefaction. 
Under  ordinary  atmospheric  pressure  the  liquid  boils  at  -  195.5°  > 
the  gas,  therefore,  can  be  liquefied  by  the  cold  obtained  by  the 
rapid  evaporation  of  liquid  oxygen  (see  p.  78). 


CHAPTER   V 

OXIDES  AND  OXY-ACIDS  OF  NITROGEN 
NITROGEN  combines  with  oxygen,  forming  five  oxides: — 

(i.)  Nitrous  oxide  (hyponitrous  anhydride)  .  N2O. 

(2.)  Nitric  oxide NO. 

(3.)  Nitrogen  trioxide  (nitrous  anhydride)    .  N2O3. 

(4.)  Nitrogen  peroxide*        ....  NO2  and  N2O4. 

(5.)  Nitrogen  pentoxide  (nitric  anhydride)    .  N2O5. 

Three  oxy-acids  of  nitrogen  are  known,  corresponding  to  the 
three  oxides,  Nos.  i,  3,  5  : — 

Hyponitrous  acid          .         .         .     H2N2O2. 
Nitrous  acid          ....     HNO2. 
Nitric  acid HNO3. 

The  most  important  of  all  these  compounds,  and  the  one  from 
which  all  the  others  are  directly  or  indirectly  obtained,  is  nitric 
acid. 

NITRIC  ACID. 

Formula,  HNC>3.     Molecular  weight— 63.02. 

History. — Nitric  acid,  or  aquafortis,  was  a  well-known  and 
valued  liquid  to  the  alchemists.  Down  to  the  time  of  Lavoisier 

*  This  name  is  usually  applied  to  this  substance  both  at  low  temperatures 
when  its  composition  is  expressed  by  the  formula  N2O4,  and  also  at  higher 
temperatures  when  the  molecules  have  dissociated  into  the  simpler  molecules 
NO2.  In  the  strictest  sense,  however,  they  may  be  regarded  as  two  oxides, 
and  it  has  been  suggested  to  name  the  one  (N2O4)  nitrogen  tetroxide,  and  the 
other  nitrogen  peroxide. 

234 


Nitric  Acid  235 

(1776)  its  true  nature  was  not  known  ;  he  showed  that  oxygen  was 
one  of  its  constituents,  but  as  to  its  other  components  he  was  un- 
certain. Its  exact  composition  was  determined  by  Cavendish. 

Modes  of  Formation. — (i.)  When  an  electric  spark  is  passed 
through  a  detonating  mixture  of  oxygen  and  hydrogen  with  which 
a  certain  quantity  of  air  or  nitrogen  is  mixed,  the  water  that  is 
produced  by  the  union  of  the  oxygen  and  hydrogen  is  found  to 
contain  nitric  acid.  This  fact  was  first  observed  by  Cavendish  in 
the  course  of  his  investigations  on  the  composition  of  water,  when, 
owing  to  the  accidental  admixture  of  air  with  the  mixed  gases, 
oxygen  and  hydrogen,  he  found  that  the  water  resulting  from  the 
union  was  sometimes  acid. 

The  direct  union  of  nitrogen  and  oxygen  may  be  brought  about 
by  allowing  a  series  of  electric  sparks  to  pass  between  platinum 
wires  in  a  confined  volume  of  air,  contained  in  a  glass  globe,  as 
shown  in  Fig.  49.  In  a  short  time  the  air  in  the  globe  will  become 
distinctly  reddish  in  colour,  owing  to  the  formation  of  nitrogen 
peroxide.  The  rapidity  of  the  formation  of  the  red  fumes  will 
be  greatly  increased  by  compressing  the  air  within  the  globe  by 
means  of  a  small  compression  pump,  as  indicated  in  the  figure. 

If  a  small  quantity  of  water  be  introduced,  and  the  contents  of 
the  globe  shaken  up,  the  red  gas  will  be  seen  to  dissolve  in  the 
water,  which  will  then  acquire  an  acid  reaction,  owing  to  the  forma- 
tion of  nitric  acid. 

Similarly,  when  a  jet  of  hydrogen  is  allowed  to  burn  in  air  to 
which  additional  oxygen  has  been  added,  considerable  quantities 
of  nitrogen  peroxide  are  formed.  The  hydrogen  may  be  burnt 
from  a  jet,  surrounded  by  a  glass  tube,  as  shown  in  Fig.  50,  into 
which  oxygen  can  be  passed  by  means  of  the  small  bent  tube  at 
the  bottom.  On  holding  a  clean  dry  cylinder  over  the  flame, 
sufficient  of  the  products  of  combustion  will  collect  in  a  few  seconds 
to  show  the  presence  of  nitrogen  peroxide. 

This  direct  union  of  atmospheric  oxygen  and  nitrogen  has 
recently  been  made  the  basis  of  a  manufacturing  process.  A 
stream  of  air  is  caused  to  pass  through  the  electric  arc  at  a  rate 
sufficiently  rapid  to  sweep  away  the  products  of  the  action  and  so 
prevent  their  dissociation.  The  nitrogen  peroxide  which  is  formed 
is  condensed  to  the  liquid  state,  and  thereby  separated  from  the 
other  gases,  by  passing  the  mixture  through  a  refrigerator. 

(2.)  Nitric  acid  is  formed  when  nitrogenous  animal  matter  under- 
goes slow  oxidation  in  the  air,  in  the  presence  of  water  and  an 


236 


Inorganic  Chemistry 


alkali,  the  nitric  acid  combining  with  the  alkali  to  form  a  nitrate. 
In  this  way  nitrates  are  found  in  the  soil,  and  from  the  soil  often 
find  their  way  into  shallow  well-waters  of  towns.  In  hot  and  rain- 
less countries  these  nitrates  are  sometimes  found  as  crystalline 
deposits  on  the  surface  of  the  soil,  as  in  Chili  and  India  (see 
Potassium  Nitrate). 

(3.)  Nitric  acid  is  prepared  by  acting  upon  potassium  nitrate 
(nitre-saltpetre)  with  sulphuric  acid.  The  nitre  is  placed  in  a  glass 
retort,  together  with  an  equal  weight  of  sulphuric  acid,  and  the 
mixture  gently  heated.  The  nitric  acid  readily  distils  over,  and 


FIG.  49. 


FIG.  50. 


may  be  collected  in  a  cooled  receiver.  The  residue  in  the  retort 
consists  of  hydrogen  potassium  sulphate — 

KN03  +  H2SO4=KHSO4  +  HNO3. 

The  acid  so  obtained  is  not  entirely  free  from  water,  and  contains 
nitrogen  peroxide  in  solution,  which  imparts  to  it  a  yellowish-red 
colour.  To  purify  it,  it  is  again  distilled  with  an  equal  volume  of 
sulphuric  acid  ;  and  the  redistilled  acid  is  deprived  of  the  last  traces 
of  dissolved  peroxide  of  nitrogen,  by  causing  a  stream  of  dry  air  to 


Nitric  Acid  237 

bubble  through  it  while  slightly  warm.     Nitric  acid  so  prepared 
may  contain  as  much  as  99.8  per  cent,  of  anhydrous  acid,  HNO3. 

(4.)  Nitric  acid  is  an  article  of  commercial  manufacture.  In  this 
process  potassium  nitrate  is  replaced  by  the  sodium  salt,  as  being 
the  cheaper  material.  The  proportion  of  acid  to  sodium  nitrate 
employed  was  formerly  arranged  in  accordance  with  the  equation  — 


It  will  be  seen  that  the  whole  of  the  hydrogen  of  the  sulphuric 
acid  is  thus  replaced  by  the  alkali  metal  derived  from  two  molecules 
of  the  nitrate,  and  that  two  molecules  of  nitric  acid  result. 

This  reaction  takes  place  in  two  stages  ;  in  the  first  we  have  — 

(i)     NaNO3  +  H2SO4  =  NaHSO4  +  HNO3. 

And  then,  as  the  temperature  is   raised,  the   hydrogen  sodium 
sulphate  reacts  upon  a  second  molecule  of  the  nitrate,  thus  — 


(2)     NaNO3  +  NaHSO4  =  Na2SO4+HNO3. 

The  temperature  necessary  to  effect  this  second  stage,  however, 
causes  the  decomposition  of  a  certain  quantity  of  the  nitric  acid  — 

2HNO3  =  H2O  +  2NO2  +  O. 

And  for  this  and  other  reasons,  most  modern  manufacturers 
work  only  to  equation  No.  I. 

The  retorts  usually  employed  for  the  manufacture  of  this  acid 
are  large  cast-iron  stills,  which  are  sometimes  lined,  either  entirely 
or  in  part,  with  fireclay,  and  which  are  built  into  a  furnace  in  such 
a  manner  as  to  allow  of  their  being  heated  as  uniformly  as  possible. 
The  charge  of  sodium  nitrate  (12  to  14  cwts.)  and  sulphuric  acid  is 
introduced,  and  the  vapours  carried  off  through  an  earthenware 
pipe  (i*,  Fig.  51),  connected  to  a  series  of  earthenware  pots,  b,  in  the 
manner  shown  in  the  figure.  The  last  of  these  jars  is  connected 
with  a  tower,  filled  with  coke,  down  which  water  is  caused  to  per- 
colate, and  any  peroxide  of*  nitrogen  which  escapes  is  thereby 
absorbed.  The  most  modern  form  of  still  is  not  cylindrical,  as 
shown  in  Fig.  51,  but  takes  the  shape  of  an  enormous  crucible 
with  a  dome-shaped  lid  ;  and  is  furnished  with  an  exit  pipe  at  the 
bottom,  from  which  the  liquid  sodium  bisulphate  is  run  off. 

Properties.—  Nitric  acid  is  a  colourless  liquid  having  a  specific 
gravity  of  1.53.  It  fumes  strongly  in  the  air,  and  has  a  peculiar 
and  choking  smell.  1  1  is  extremely  hygroscopic,  absorbing  moisture 


Inorganic  Chemistry 

from  the  air  with  great  readiness.  Nitric  acid  is  an  intensely 
corrosive  liquid  :  the  strongest  acid,  when  brought  in  contact  with 
the  skin,  causes  painful  wounds,  while  in  more  dilute  conditions 


it  stains  the  skin  and  other  organic  materials  a  bright  yellow 
colour.  A  quantity  of  strong  nitric  acid  thrown  upon  sawdust 
causes  it  to  burst  into  flame.  When  nitric  acid  is  distilled  it  first 


Nitric  Acid  239 

begins  to  boil  at  86°,  at  the  same  time  it  is  partially  decomposed 
into  water,  nitrogen  peroxide,  and  oxygen  ;  the  distillate,  therefore, 
gradually  becomes  weaker,  and  the  boiling-point  gradually  rises. 
This  continues  until  a  certain  point  is  reached,  when  both  the 
temperature  of  the  boiling  liquid  and  the  strength  of  the  distillate 
remain  constant.  If,  on  the  other  hand,  a  weak  acid  be  distilled, 
the  distillate  gradually  increases  in  strength,  until,  when  the  same 
point  is  reached,  the  boiling  liquid  has  again  the  same  temperature. 

This  constant  boiling-point  is  120.5°,  and  ^e  distillate  which 
comes  over  at  that  temperature  contains  68  per  cent,  of  HNO3. 
Whatever  the  strength  of  the  acid,  therefore,  on  being  boiled  it 
loses  either  nitric  acid  or  water  until  the  strength  reaches  68  per 
cent,  and  this  liquid  boils  at  120°  C.  The  specific  gravity  of  this 
acid  at  15°  is  1.414.  It  was  formerly  supposed  that  the  acid  of  this 
strength  constituted  a  definite  hydrate,  but  Roscoe  has  shown  that 
the  strength  of  the  acid  is  purely  a  function  of  the  pressure,  for  by 
varying  the  pressure  under  which  the  distillation  is  conducted, 
acids  of  various  compositions  can  be  caused  to  distil  at  a  constant 
temperature.  Mixed  liquids  of  this  nature  are  known  as  constant- 
boiling  mixtures,  and  are  strictly  analogous  to  constant-freezing 
mixtures  (page  155). 

When  nitric  acid  is  mixed  with  water  there  is  a  rise  in  tempera- 
ture and  a  contraction  in  volume,  the  maximum  effect  being  pro- 
duced when  the  mixture  is  made  in  the  proportion  of  three  molecules 
of  water  with  one  molecule  of  acid. 

Nitric  acid  is  a  powerful  oxidising  agent,  on  account  of  the  readi- 
ness with  which  it  parts  with  oxygen.  Elements  such  as  sulphur 
and  phosphorus  are  oxidised  into  sulphuric  and  phosphoric  acids  ; 
arsenious  oxide  into  arsenic  acid  ;  and  many  protosalts  are  con- 
verted into  persalts.  It  attacks  a  large  number  of  metals,  forming 
in  many  cases  the  nitrate.  Its  action  upon  metals  is  often  of 
a  complicated  nature,  and  depends  not  only  upon  the  particular 
metal,  but  also  upon  the  strength  of  the  acid,  the  temperature, 
and  the  presence  of  the  saline  products  of  the  reaction  ;  thus, 
when  nitric  acid  acts  upon  copper,  the  following  reaction  takes 
place  — 

3Cu+8HNO3=3Cu(NO3) 


It  is  found,  however,  that  as  the  amount  of  copper  nitrate  accu- 
mulates, the  nitric  oxide  which  is  evolved  is  mixed  more  and  more 
largely  with  nitrous  oxide,  N?O,  and  even  with  nitrogen. 


240  Inorganic  Chemistry 

Again,  when  dilute  nitric  acid  acts  upon  zinc,  nitrous  oxide  is 
produced,  according  to  the  following  equation  — 

4Zn  +  10HNO3  =  4Zn(NO3)2  +  5H2O  +  N2O. 

When,  however,  strong  nitric  acid  i.s  employed,  ammonia  is  formed, 
which  combines  with  the  excess  of  acid  — 

4Zn  +  9HNO3=4Zn(NO3)2  +  3H2O  +  NH3. 

In  some  cases,  as  with  copper  and  silver,  the  presence  of  nitrous 
acid  (either  as  an  impurity  in  the  nitric  acid,  or  as  a  first  product 
of  its  attack  upon  the  metal)  is  believed  to  be  a  necessary  condition 
of  the  action. 

Owing  to  the  strong  oxidising  properties  of  nitric  acid,  hydro- 
gen is  rarely  isolated  by  the  action  of  metals  upon  this  acid,  the 
hydrogen  which  is  displaced  from  the  acid  being  converted  into 
water.  With  magnesium,  however,  free  hydrogen  is  evolved. 

The  chief  reactions  of  nitric  acid  may  be  broadly  divided  into 
three  classes  :  — 

(i.)  With  metallic  oxides  its  behaviour  is  in  common  with  other 
acids.  It  exchanges  its  hydrogen  for  an  equivalent  quantity  of  the 
metal,  forming  a  nitrate,  with  the  elimination  of  water,  e.g.  — 


(2.)  Reactions  in  which  it  acts  as  an  oxidising  agent  ;  as  an 
example,  its  action  upon  iodine,  which  is  converted  into  iodic  acid, 
may  be  cited  — 

I  +  3HN03=HIO3+H2O  +  NO  +  2NO2. 

'  (3.)  Actions  in  which  hydrogen  in  an  organic  compound  is 
replaced  by  the  elements  NO2,  with  the  elimination  of  H2O,  no 
gas  being  evolved.  The  conversion  of  cotton-wool,  or  cellulose, 
C]2H20O10,  into  gun-cotton,  or  nitro-cellulose,  CI2H14O10(NO2)6,  is 
an  illustration  of  this  class  of  reactions  — 

C12H20010-f6HN03  =  6H20  +  C12H14010(N02)6. 

Nitric  acid  is  without  action  upon  the  so-called  noble  metals, 
gold  and  platinum. 

Commercial  nitric  acid,  which  is  of  a  reddish  colour,  is  liable 
to  contain  many  impurities  :  chlorine  and  iodic  acid,  derived  from 
the  Chili  saltpetre  ;  iron,  sulphuric  acid,  and  sodium  sulphate, 


Nitrogen  Pentoxide  241 

carried  mechanically  over  from  the  retorts  ;  and  nitrogen  peroxide, 
from  the  decomposition  of  the  acid.  From  these  it  is  purified  by 
redistillation. 

Nitric  acid  is  a  monobasic  acid  ;  the  salts  of  which,  known  as 
the  nitrates,  are  for  the  most  part  readily  soluble  in  water,  and 
crystallise  in  well-defined  forms.  They  are  all  decomposed  at  a 
high  temperature,  evolving  oxygen  and  nitrogen  peroxide,  or  oxy- 
gen and  nitrogen,  leaving  an  oxide  of  the  metal. 

The  presence  of  a  nitrate  in  solution  is  easily  recognised  by  the 
following  characteristic  test.  A  solution  of  ferrous  sulphate  is  first 
added  to  the  solution  containing  the  nitrate,  and  concentrated  sul- 
phuric acid  is  then  cautiously  poured  down  the  side  of  the  test- 
tube,  held  in  a  sloping  position,  so  as  to  fall  to  the  bottom  without 
mixing  with  the  solution.  The  sulphuric  acid  acting  upon  the 
nitrate  liberates  nitric  acid  ;  this  is  reduced  by  the  ferrous  sulphate 
to  nitric  oxide,  which,  dissolving  in  the  ferrous  sulphate,  forms  a 
brown-coloured  solution  at  the  point  where  the  two  layers  of  liquid 
meet  (see  Nitric  Oxide). 

When  nitric  acid  is  added  to  hydrochloric  acid,  a  mixture  is 
obtained  which  is  known  by  the  name  of  aqua  regia.  This  name 
•was  applied  to  it  by  the  alchemists  on  account  of  its  power  of  dis- 
solving gold.  Aqua  regia  is  used  in  the  laboratory  for  dissolving 
gold,  platinum,  and  certain  ores.  Its  solvent  power  depends  upor 
the  free  chlorine  which  is  evolved  from  the  mixture — 


NITROGEN  PENTOXIDE  (Nitric  Anhydride). 
Formula,  N2O3.     Molecular  weight  =108.02. 

Modes  Of  Formation.—  (  i.)  By  withdrawing  from  nitric  acid 
the  elements  of  water,  by  means  of  phosphorus  pentoxide  — 


For  this  purpose  the  strongest  nitric  acid  is  cautiously  added  to 
phosphorus  pentoxide  in  a  cooled  retort,  in  the  proportion  de- 
manded by  the  equation  ;  the  mixture  being  made  as  far  as  possible 
without  rise  of  temperature.  The  pasty  mass  is  then  gently  heated, 
when  the  nitrogen  pentoxide  distils  over,  and,  if  collected  in  a  well- 
cooled  receiver,  at  once  crystallises. 

(2.)  The  method  adopted  by  Deville,  who  discovered  this  com- 

Q 


242  Inorganic  Chemistry 

pound  (1849),  was  by  passing  dry  chlorine  over  dry  silver  nitrate 
contained  in  a  U-tube,  which  was  kept  at  the  desired  temperature 
by  being  immersed  in  a  water-bath.  The  following  equation  ex- 
presses the  final  result  of  the  action  — 


Properties.  —  Nitrogen  pentoxide  is  a  white  solid  substance, 
crystallising  in  brilliant  prismatic  crystals,  which  melt  at  30° 
with  partial  decomposition.  Between  45°  and  50°  it  undergoes 
rapid  decomposition,  evolving  brown  fumes.  It  is  a  very  unstable 
compound  ;  when  suddenly  heated  it  decomposes  with  explosive 
violence,  and  even  at  ordinary  temperatures  decomposition  slowly 
takes  place.  It  absorbs  moisture  rapidly,  and  when  thrown  into 
water  it  dissolves  with  the  evolution  of  great  heat  — 

=  2HNO3. 


When  nitrogen  pentoxide  is  gradually  mixed  with  nitric  acid,  a 
compound  is  formed  having  the  composition  2N2O5,H2O  ;  which 
separates,  on  cooling,  as  a  definite  crystalline  hydrate 

NITROGEN  PEROXIDE. 

Formula,  NO2  and  N2O4.     Molecular  weight  =46.01  and  92.02. 
Density  =23.00  and  46.01. 

Modes  Of  Formation.—  (  i.)  This  compound  may  be  prepared 
by  mixing  one  volume  of  oxygen  with  two  volumes  of  nitric  oxide, 
and  passing  the  red  gas  so  obtained  through  a  tube  surrounded 
by  a  freezing-mixture  — 

=  2NO. 


(2.)  The  nitrates  of  certain  metals,  when  heated,  are  decomposed 
into  nitrogen  peroxide,  oxygen,  and  an  oxide  of  the  metal  ;  thus, 
if  dry  lead  nitrate  be  heated  in  a  retort  and  the  gaseous  products 
of  decomposition  are  conducted  into  a  U-tube  placed  in  a  freezing- 
mixture,  the  nitrogen  peroxide  collects  in  the  tube  — 


(3.)  When  arsenious  oxide  is  gently  warmed  with  nitric  acid,  a 
mixture  of  nitric  oxide,  NO,  and  peroxide,  NO2,  is  evolved,  and  if 
this  gaseous  mixture  be  passed  through  a  cooled  tube,  it  condenses 


Nitrogen  Peroxide  243 

to  a  blue  liquid.  On  passing  a  stream  of  oxygen  through  this 
liquid  it  loses  its  blue  colour,  and  is  converted  into  a  yellowish 
liquid  which  consists  of  nitrogen  peroxide. 

Properties. — At  low  temperatures  nitrogen  peroxide  is  a  colour- 
less crystalline  compound.  It  melts  at  -  9°,  but  requires  a  tem- 
perature as  low  as  —  30°  to  solidify  it.  At  a  temperature  slightly 
above  its  melting-point  the  liquid  begins  to  acquire  a  pale  yellowish 
tint,  which  rapidly  deepens  until  at  the  ordinary  temperature  it  is 
a  full  orange  colour.  The  liquid  boils  at  22°,  and  gives  a  vapour 
having  a  reddish-brown  colour.  The  colour  of  the  vapour  also 
becomes  deeper  as  its  temperature  is  raised,  until  at  40°  it  is  a 
dark  chocolate  brown,  and  almost  opaque.  On  allowing  the 
vapour  to  cool  the  reverse  changes  take  place.  This  change  of 
colour,  as  the  temperature  rises,  is  accompanied  by  a  steady 
change  in  the  density  of  the  gas,  as  will  be  seen  from  the  table  : — 

Temperature.  Density.  ^^^,1 

26.7°  38.3  20.00 

6o.2°  30.1  50.04 

100.1°  24.3  79.23 

135.0°  23.1  98.96 

140.0°  23.00  loo.oo 

The  density  required  by  the  formula  N2O4  is  46.04,  while  that 
demanded  by  the  formula  NO2  is  23.02  ;  hence  as  the  temperature 
rises  a  process  of  dissociation  goes  on  in  which  N2O4  molecules 
are  broken  down  into  molecules  of  the  simpler  composition. 
At  140°  this  process  is  complete,  and  the  gas  is  entirely  re- 
solved into  NO2.  It  is  believed  that  at  low  temperatures  nitrogen 
peroxide  has  the  composition  represented  by  the  formula  N2O4, 
but  that  dissociation  begins  to  take  place  even  during  the  state  of 
liquidity,  as  indicated  by  the  gradual  change  of  colour  ;  and  there- 
fore at  temperatures  between  the  boiling-point  of  the  liquid,  viz., 
22°,  and  140°,  the  gas  consists  of  mixtures  of  molecules  of  NO2  and 
N2O4.  The  calculated  percentage  of  NO2  molecules,  which  the 
gas  contains  at  the  temperatures  at  which  the  above  densities  are 
taken,  are  given  in  the  third  column. 

Nitrogen  peroxide  is  decomposed  by  water.  At  low  tempera- 
tures,  and  with  small  quantities  of  water,  nitric  and  nitrous  acids 
are  the  products  of  the  action,  thus — 

N2O4+H2O  = 


244  Inorganic  Chemistry 

At  the  ordinary  temperature,  and  in  the  presence  of  an  excess  of 
water,  the  following  reaction  takes  place  — 

3NO2  +  H2O  =  2HNO3  +  NO. 

Gaseous  nitrogen  peroxide  is  incapable  of  supporting  the  com- 
bustion of  a  taper.  Phosphorus,  when  strongly  burning  and 
plunged  into  the  gas,  continues  its  combustion  with  brilliancy, 
the  temperature  of  the  burning  phosphorus  being  sufficiently  high 
to  effect  the  decomposition  of  the  gas.  Nitrogen  peroxide  is  a 
suffocating  and  highly  poisonous  gas,  and  even  when  largely 
diluted  with  air  rapidly  produces  headache  and  sickness. 

Nitrogen  peroxide  unites  directly  with  certain  metals,  giving  rise  to  a  re- 
markable series  of  compounds,  to  which  the  name  nitro-metals,  or  metallic 
nitroxyls,  may  be  given  (Sabatier  and  Senderens).*  Thus,  when  the  vapour 
of  nitrogen  peroxide  is  passed  over  metallic  copper  (obtained  by  the  reduction 
of  copper  oxide  in  a  stream  of  hydrogen),  the  gas  is  rapidly  absorbed  by  the 
metal  with  considerable  rise  of  temperature,  and  a  solid  brown  compound  is 
formed.  This  substance  is  the  copper-nitroxyl,  and  its  composition  is  ex- 
pressed by  the  formula  Cu.2NO2. 

Copper-nitroxyl  is  a  fairly  stable  compound,  and  is  unacted  upon  by  dry  air. 
It  is  decomposed  by  water  and  by  nitric  acid,  hence  in  its  preparation  care 
must  be  taken  to  free  the  nitrogen  peroxide  from  these  substances. 

At  a  temperature  of  about  90°  copper-nitroxyl  is  decomposed  into  copper 
and  nitrogen  peroxide.  If,  therefore,  a  quantity  of  the  compound  be  sealed 
up  in  a  bent  glass  tube,  and  the  empty  limb  of  the  tube  be  immersed  in  a 
freezing-mixture  while  the  compound  is  gently  warmed,  the  nitrogen  peroxide 
which  is  evolved  will  be  condensed  in  the  cold  portion  of  the  tube. 

Similar  compounds  are  formed  with  the  metals  cobalt,  nickel,  and  iron. 

Nitrous  Acid,  HNO2.—  This  substance  is  not  known  in  the  pure 
state.  Even  in  dilute  aqueous  solution  it  rapidly  decomposes  into 
nitric  acid,  nitric  oxide,  and  water  — 


The  solution  of  this  acid  sometimes  acts  as  a  reducing  agent. 
taking  up  oxygen  from  such  highly  oxidised  compounds  as  per- 
manganates or  chromates  and  passing  into  nitric  acid  — 


HNO2  +  O  =  HNO3. 


Under  other  conditions  it  exerts  an  oxidising  action,  as  when  it 
bleaches  indigo,  or  liberates  iodine  from  potassium  iodide,  being 

*  Bulletin  de  la  Socittt  Chimique,  September  1893. 


Nitrous  Acid  245 

itself  reduced  to  nitric  oxide  and  water,  with  the  elimination  of 
oxygen  — 

2HNO2  =  2NO  +  H2O  +  O. 

The  salts  of  nitrous  acids,  viz.,  the  nitrites,  are  stable  compounds. 
The  alkali  nitrites  may  be  prepared  by  carefully  heating  the  nitrates 
above  their  fusion  points  — 

KNO3  =  KNO2  +  O. 
At  a  higher  temperature  the  nitrite  is  also  decomposed. 

On  the  manufacturing  scale  the  nitrate  is  reduced  by  fusion  with 
metallic  lead  — 


Nitrites  are  decomposed  by  dilute  acids  evolving  brown  vapours, 
and  in  this  way  are  at  once  distinguished  from  nitrates. 

Nitrogen  Trioxide,  N2O3.—  This  compound  is  obtained  by 
passing  a  mixture  of  nitric  oxide  and  nitric  peroxide  through  a 
tube  cooled  to  about  -20°,  when  it  condenses  in  the  form  of  a 
bluish  green  liquid. 

Nitrogen  trioxide  so  obtained  is  stable  only  at  low  temperatures. 
When  the  temperature  is  allowed  to  rise  the  liquid  itself  undergoes 
dissociation  into  its  two  generators,  and  as  the  more  volatile  oxide, 
NO,  escapes,  the  green  colour  of  the  liquid  gradually  disappears, 
leaving  a  yellow  liquid  which  finally  passes  off  as  gaseous  NO2. 
The  reaction  — 

.NO  +  NO2;±N2O3 

is  therefore  a  reversible  reaction  in  which  the  change  as  read 
from  right  to  left  is  practically  complete  at  ordinary  temperatures. 
There  has  long  been  considerable  doubt  as  to  whether  N2O3 
had  any  existence  as  a  gas.  All  the  ordinary  chemical  reactions 
in  which  this  compound  was  at  one  time  thought  to  be  evolved  as 
gas,  were  found  in  reality  to  yield  mixtures  of  NO  and  NO2  ;  thus 
in  the  case  of  the  action  of  nitric  acid  upon  arsenious  oxide  — 


It  has  now,  however,  been  shown  by  Baker  (Jour.  Chem.  Soc., 
Nov.  1907)  that  if  the  liquid  N2O3  is  rendered  absolutely  dry  by 
prolonged  exposure  to  phosphoric  oxide,  it  then  vaporises  com- 
pletely without  dissociating,  yielding  a  vapour  whose  density  was 
in  no  case  below  38  (that  demanded  by  the  formula  N2O3),  and 
which  in  many  experiments  was  considerably  above  this  figure.  This 
latter  fact  would  seem  to  indicate  that  under  these  conditions  the 
gas  may  to  some  extent  be  polymerised,  probably  into  N4O6,  cor- 
responding to  the  analogous  oxides  of  phosphorus  and  arsenic. 


246  Inorganic  Chemistry 

NITRIC   OXIDE. 

Formula,  NO.     Molecular  weight  =30.  04.     Density  =  15.  02. 

History.  —  Nitric  oxide  was  first  obtained  by  Van  Helmont. 
Priestley,  however,  was  the  first  to  investigate  this  gas,  which  he 
termed  nitrous  air,  and  which  was  employed  by  him  in  his  analysis 
of  air. 

Modes  of  Formation.—  (  i.)  This  gas  is  obtained  by  the  action 
of  nitric  acid  of  specific  gravity  1.2  upon  copper  or  mercury.  In 
practice  copper  is  always  employed.*  The  action  may  be  repre- 
sented thus  — 

3Cu  +  8HNO3  =  3Cu(NO3)2  +  4H2O  +  2NO. 

The  gas  obtained  by  this  method  is  always  liable  to  contain 
nitrous  oxide  and  even  free  nitrogen  ;  the  amount  of  these  im- 
purities rapidly  increasing  if  the  temperature  be  allowed  to  rise, 
and  still  more  so  as  the  amount  of  copper  nitrate  in  solution 
increases. 

(2.)  Pure  nitric  oxide  is  readily  obtained  by  the  action  of  nitric 
acid  upon  ferrous  sulphate.  The  reaction  is  best  applied  by  gene- 
rating the  nitric  acid  from  potassium  nitrate  and  sulphuric  acid  in 
the  presence  of  ferrous  sulphate.  A  mixture  of  the  two  salts,  in 
the  proportion  of  about  one  part  of  nitre  to  four  of  ferrous  sulphate, 
is  introduced  into  a  flask,  with  a  small  quantity  of  water.  Strong 
sulphuric  acid  is  dropped  upon  the  mixture  by  means  of  a  drop- 
ping funnel,  and  the  mixture  gently  warmed,  when  a  steady  stream 
of  pure  nitric  oxide  is  evolved  — 

2KNO3  +  5H2SO4  +  6FeSO4=2HKSO4+3Fe2(SO4)3  +  4H2O  +  2NO. 

A  precisely  similar  result  may  be  obtained  by  the  reduction  of 
potassium  nitrate  by  means  of  ferrous  chloride  in  the  presence  of 
hydrochloric  acid,  thus  — 


Properties.  —  Nitric  oxide  is  a  colourless  gas,  having  a  specific 
gravity  of  1.039.  When  brought  into  the  air,  it  combines  with  the 
atmospheric  oxygen,  forming  red-brown  vapours,  consisting  of 

*  Experiment  314,  "  Chemical  Lecture  Experiments,"  new  ed. 


Nitric  Oxide 


247 


nitrogen  peroxide,  the  combination  being  attended  with  a  rise 
of  temperature.  The  formation  of  these  red  fumes  in  contact 
with  oxygen  is  characteristic  of  this  gas,  thereby  distinguishing 
it  from  all  other  gases.  This  property  of  nitric  oxide  renders 
it  impossible  to  ascertain  whether  this  gas  has  any  smell,  or 
is  possessed  of  any  toxicological  action.  Nitric  oxide  is  only 
very  sparingly  soluble  in  water.  It  is  the  most  stable  of  all  the 
oxides  of  nitrogen,  being  able  to  stand  a  dull  red  heat  without 
decomposition.  It  is  not  a  supporter  of  combustion.  A  lighted 
taper,  or  a  burning  piece  of  sulphur,  when  introduced  into  the  gas, 
are  extinguished.  If  the  temperature  of  the  burning  substance  is 
sufficiently  high  to  decompose  the  gas,  com- 
bustion then  continues  at  the  expense  of  the 
liberated  oxygen  :  thus,  if  a  piece  of  phos- 
phorus, which  is  freely  burning  in  the  air,  be 
plunged  into  this  gas,  it  continues  its  com- 
bustion with  great  brilliancy  ;  if,  however, 
the  phosphorus  be  only  feebly  burning  when 
thrust  into  the  gas,  it  is  at  once  extinguished. 
A  mixture  of  carbon  disulphide  vapour  and 
nitric  oxide,  obtained  by  allowing  a  few  drops 
of  the  liquid  to  fall  into  a  cylinder  of  the  gas, 
burns,  when  inflamed,  with  an  intensely  vivid 
bluish  flame,  which  is  especially  rich  in  the 
violet  or  actinic  rays,  and  has  on  this  ac- 
count been  sometimes  employed  by  photo- 
graphers to  illuminate  dark  interiors.  Nitric 

oxide  is  soluble  in  a  solution  of  ferrous  sulphate,  forming  a  dark- 
brown  solution,  containing  an  unstable  compound  of  ferrous  sul- 
phate and  nitric  oxide,  2FeSO4,NO.  This  compound  is  readily 
decomposed  by  heat,  nitric  oxide  being  evolved.  By  means  of  this 
reaction,  nitric  oxide  may  be  separated  from  other  gases.  Nitric 
oxide  is  a  difficultly  liquefiable  gas,  its  critical  temperature  being 
~93-5°:  at  this  temperature  a  pressure  of  71.2  atmospheres  is 
required  to  liquefy  it. 

The  composition  of  nitric  oxide  may  be  proved  by  heating  a 
spiral  of  iron  wire  by  means  of  an  electric  current  in  a  measured 
volume  of  the  gas  (as  shown  in  Fig.  52).*  As  the  metal  becomes 
red  hot  the  gas  is  gradually  decomposed  and  the  oxygen  combines 


FIG.  52. 


*  No.  321,  "Chemical  Lecture  Experiments.1 


248  Inorganic  Chemistry 

with  the   iron  to  form  ferric  oxide.     The  residual  nitrogen  will 
be  found  to  occupy  one-half  the  original  volume. 

Two  vols.  of  nitric  oxide,  weighing    30.04 
Contain  i  vol.  of  nitrogen,  weighing  14.04 

1 6. 00= weight  of  i  vol.  of  oxygen. 

Therefore  we  learn  that  two  volumes  of  nitric  oxide  consist  of 
one  volume  of  nitrogen  and  one  volume  of  oxygen  united  without 
condensation. 


NITROUS  OXIDE  (Hyponitrous  anhydride,  Laughing  gas}. 
Formula,  N2O.     Molecular  weight =44. 08.     Density =22. 04. 

History. — This  gas  was  discovered  by  Priestley,  and  called  by 
him  dephlogisticated  nitrous  air. 

Modes  of  Formation. — (i.)  Nitrous  oxide  is  formed  by  the 
reduction  of  nitric  acid  by  certain  metals,  as  zinc  or  copper,  under 
special  conditions  (see  Nitric  Acid).  These  reactions,  however, 
are  never  made  use  of  for  the  preparation  of  the  gas  for  experi- 
mental purposes. 

(2.)  The  most  convenient  method  for  obtaining  this  compound 
is  by  the  decomposition  of  ammonium  nitrate.  A  quantity  of  the 
dry  salt  is  gently  heated  in  a  flask  fitted  with  a  cork  and  delivery- 
tube.  The  salt  rapidly  melts  and  splits  up  into  nitrous  oxide  and 
water — 

NH4NO3  =  2H2O  +  N2O. 

The  heat  should  be  carefully  regulated,  or  the  decomposition  is 
liable  to  become  violent,  in  which  case  nitric  oxide  is  also  evolved. 
Nitrous  oxide  being  rather  soluble  in  cold  water,  the  gas  should 
be  collected  either  over  mercury  or  over  hot  water. 

When  the  gas  is  to  be  used  for  anaesthetic  purposes,  it  should  be  purified 
by  being  passed  first  through  a  solution  of  ferrous  sulphate  to  absorb  any  nitric 
oxide,  and  afterwards  through  caustic  soda,  to  remove  any  chlorine  which  may 
have  been  derived  from  the  presence  of  ammonium  chloride  in  the  nitrate. 

Properties. — Nitrous  oxide  is  a  colourless  gas,  having  a  faint 
and  not  unpleasant  smell,  and  a  peculiar  sweetish  taste.  Its 
specific  gravity  is  1.52.  The  gas  is  somewhat  soluble  in  water,  its 
coefficient  of  absorption  at  o°  being  1.3052.  The  solubility  rapidly 


Nitrous  Oxide  249 

decreases  as  the  temperature  rises,  as  will  be  seen  by  the  follow- 
ing table  (Carius)  : — 

i  c.c.  Water  at  c.c.  N2O  at  o°  C. 

760  mm.  Dissolves  and  760  mm. 

At    o° 1.3052 

„    10°  ,  0.9196 

,,20°                        .            .             .            .      0.6700 
,,25° C.5962 

The  loss  of  gas  during  its  collection  over  water  in  the  pneumatic 
trough,  arising  from  its  solubility  in  that  liquid,  is  therefore  greatly 
lessened  by  using  warm  water.  Nitrous  oxide  is  much  more 
readily  decomposed  than  nitric  oxide  ;  a  red-hot  splint  of  wood  is 
instantly  rekindled,  and  bursts  into  flame  when  plunged  into  the 
gas.  Phosphorus  burns  in  it  with  a  brilliancy  scarcely  perceptibly 
less  dazzling  than  in  pure  oxygen.  If  a  piece  of  sulphur  which  is 
only  feebly  burning  be  thrust  into  a  jar  of  this  gas,  the  sulphur  is 
extinguished,  the  temperature  of  the  flame  not  being  sufficiently 
high  to  decompose  the  gas.  When,  however,  the  sulphur  is 
allowed  to  get  into  active  combustion  before  being  placed  in  the 
gas,  the  combustion  continues  with  greatly  increased  brilliancy. 
In  all  cases  of  combustion  in  nitrous  oxide,  the  combustion  is 
simply  the  union  of  the  burning  body  with  oxygen,  the  nitrogen 
being  eliminated.  From  its  behaviour  towards  combustibles, 
nitrous  oxide  might  readily  be  mistaken  for  oxygen  ;  it  can,  how- 
ever, be  easily  distinguished  from  that  gas  by  the  fact  that  when 
added  to  nitric  oxide  it  does  not  produce  red  vapours,  whereas 
when  oxygen  is  mixed  with  nitric  oxide  these  coloured  fumes  are 
instantly  formed. 

When  equal  volumes  of  nitrous  oxide  and  hydrogen  are  mixed 
in  a  eudiometer,  and  an  electric  spark  passed  through  the  mix- 
ture, the  gases  combine  with  explosion,  water  being  produced  and 
nitrogen  set  free  ;  the  volume  of  nitrogen  so  resulting  being  equal 
to  that  of  the  nitrous  oxide  employed.  This  compound,  therefore, 
contains  its  own  volume  of  nitrogen,  and  half  its  own  volume  of 
oxygen.  Nitrous  oxide,  when  inhaled,  exerts  a  remarkable  action 
upon  the  animal  organism.  This  fact  was  first  observed  by  Davy. 
If  breathed  for  a  short  time,  the  gas  induces  a  condition  of  hysterical 
excitement,  often  accompanied  by  boisterous  laughter,  hence  the 
name  laughing  gas.  If  the  inhalation  be  continued,  this  is  followed 
by  a  condition  of  complete  insensibility,  and  ultimately  by  death. 


250  Inorganic  Chemistry 

On  account  of  the  ease  with  which  the  state  of  insensibility  can  be 
brought  about,  this  gas  is  extensively  employed  as  an  anaesthetic, 
especially  in  dentistry. 

Nitrous  oxide  is  a  gas  which  is  moderately  easily  liquefied  ;  at 
o°  C.  a  pressure  of  thirty  atmospheres  is  required  to  effect  its 
liquefaction. 

Liquid  nitrous  oxide  is  colourless  and  mobile  ;  it  boils  at  —  89.8°, 
and  when  dropped  upon  the  skin  produces  painful  blisters. 
When  thrown  upon  water,  a  quantity  of  the  water  is  at  once  con- 
verted into  ice  ;  mercury  poured  into  a  tube  containing  a  small 
quantity  of  the  liquid  is  instantly  frozen.  An  ignited  fragment  of 
charcoal  thrown  upon  the  liquid  floats  upon  the  surface,  at  the 
same  time  burning  with  brilliancy.  If  the  liquid  be  mixed  with 
carbon  disulphide,  and  placed  in  vacuo,  the  temperature  falls 
to  -  140°.  By  strongly  cooling  the  liquid,  contained  in  a  sealed 
tube,  Faraday  succeeded  in  solidifying  it  ;  this  may  also  be 
effected  by  the  rapid  evaporation  of  the  liquid.  The  solid  melts 
at  -  102.7°,  and  if  placed  upon  the  hand  causes  a  painful  blister  ; 
in  this  respect  it  differs  from  solid  carbon  dioxide,  which  gasifies 
without  previous  liquefaction. 

Hyponitrous  Acid,  H2N2O2  or  N2(HO)2. 

When  a  solution  of  potassium  nitrate  or  nitrite  is  acted  upon  by  sodium 
amalgam  the  salt  is  reduced  by  the  nascent  hydrogen  (evolved  by  the  action 
of  the  amalgam  upon  the  water)  with  the  formation  of  potassium  hyponitrite  — 

=2H2O+K2N2O2         or 


The.  solution,  which  is  alkaline,  is  neutralised  with  acetic  acid,  and  silver 
nitrate  added,  which  causes  the  precipitation  of  yellow  insoluble  silver 
hyponitrite,  Ag2N2O2. 

The  acid  itself  is  obtained  from  the  silver  salt  by  the  action  of  dilute  hydro- 
chloric acid.  By  employing  an  ethereal  solution  of  hydrogen  chloride  and 
evaporating  the  solution,  the  hyponitrous  acid  is  obtained  in  the  form  of  white 
deliquescent  crystals.  The  pure  acid  is  very  unstable,  exploding  spontaneously 
even  below  o°  C.  The  aqueous  solution  when  gently  warmed  is  quickly 
broken  up  into  water  and  nitrous  oxide. 

Formerly  nitrous  oxide  was  regarded  as  containing  nitrogen  in  the  mono- 
valent  condition,  and  its  constitution  was  represented  by  the  formula 
N  —  O  —  N.  It  is  more  probable,  however,  that  the  molecule  contains  the 
two  nitrogen  atoms  doubly  linked,  and  that  its  formation  by  the  decomposition 
of  hyponitrous  acid  is  expressed  by  the  equation  — 

N  :  N 


Nit  rosy  I  Chloride  2$  I 

Nitrosyl  Chloride,*  NOC1.  —  This  compound  maybe  obtained  by  the  direct 
combination  of  nitric  oxide  with  chlorine  — 

2NO  +  C12=2NOC1. 

It  is  also  formed  by  the  action  of  phosphorus  pentachloride  upon  potassium 
nitrite,  thus  — 

PC15+KN02  =  NOC1  +  POC13+KC1. 

Nitrosyl  chloride  is  formed  together  with  chlorine  when  a  mixture  of  nitric 
and  hydrochloric  acids  is  gently  heated  — 

HNO3+3HC1=NOC1  +  C12+2H20. 

Nitrosyl  chloride  is  also  readily  prepared  by  the  action  of  nitrosyl  hydrogen 
sulphate  upon  dry  sodium  chloride,  thus  — 

(NO)HSO4  +  NaCl=  NOC1  +  NaHSO4. 

Properties.  —  Nitrosyl  chloride  is  an  orange-yellow  gas,  which  easily  con- 
denses when  passed  through  a  tube  immersed  in  a  freezing-mixture,  to  an 
orange-yellow  liquid,  which  boils  at  about  -  8°.  It  is  decomposed  by  water 
into  nitrous  acid  and  hydrochloric  acid— 

NOCl  +  H20  =  HNOo+HCl. 

In  a  similar  manner  it  is  decomposed  by  metallic  oxides  and  hydroxides, 
thus—  . 

=  KNOo+  KC1  +  H2O. 


Nitrosyl  chloride  has  no  action  upon  gold  and  platinum,  but  it  attacks 
mercury  with  the  formation  of  mercurous  chloride  and  the  liberation  of  nitric 
oxide— 

2NOCl  +  2Hg=Hg2Cl2+2NO. 

Nitrogen  Oxyfluorides.  —  Two  of  these  compounds  have  been  obtained. 
Nitroxyl  fluoride,  NO2F,  is  obtained  by  the  direct  union  of  fluorine  and  nitric 
oxide  at  the  temperature  of  liquid  oxygen  (Moissan).  It  is  a  colourless  gas 
to  a  liquid  which  boils  at  —63.5.  Water  decomposes  it  into  nitric  and 
hydrofluoric  acid  — 

2O  =  HNO3+HF. 


Nitrosyl  fluoride,  NOF,  is  produced  by  the  action  of  nitrosyl  chloride  upon 
silver  fluoride  (Ruff  and  Stauber)— 

NOCl  +  AgF  =  NOF  +  AgCl. 
*  Tilden  has  shown  that  this  is  the  only  oxy-chloride  of  nitrogen  that  exists. 


CHAPTER    VI 
THE    ATMOSPHERE 

THE  atmosphere  is  the  name  applied  to  the  gaseous  mixture 
which  envelops  the  earth,  and  which  is  commonly  called  the  air. 
The  older  chemists  used  the  word  air  much  as  in  modern  times 
the  word  gas  is  employed  ;  thus  they  spoke  of  inflammable  air, 
dephlogisticated  air,  alkaline  air,  and  so  on. 

The  air  consists  of  a  mixture  of  gases,  the  two  chief  ingredients 
being  nitrogen  and  oxygen.  Lavoisier  was  the  first  to  clearly 
prove  that  oxygen  was  a  constituent  of  the  air,  although  Robert 
Boyle  and  others  before  him  had  shown  that  air  was  absorbed  by 
metals  in  the  process  of  forming  a  calx,  and  that  the  metal  gained 
weight  as  the  calx  formed.  When  the  fact  that  the  air  was  com- 
posed of  oxygen  and  nitrogen  became  established,  various  de- 
vices were  adopted  to  determine  the  proportion  of  oxygen  in  it. 
Priestley's  method  was  by  means  of  nitric  oxide.  It  depended 
upon  the  fact  that  when  nitric  oxide  is  mixed  with  air  it  combines 
with  the  oxygen,  forming  brown  fumes  which  dissolve  in  the  water. 
A  contraction  in  volume  therefore  takes  place,  from  which  the 
volume  of  oxygen  may  be  calculated.  This  method  yielded  results 
which  seemed  to  show  that  there  was  considerable  variation  in  the 
proportion  of  oxygen  present  in  different  samples  of  air,  and  the 
idea  arose  that  the  wholesomeness  or  goodness  of  the  air  was 
dependent  upon  the  quantity  of  oxygen  which  it  contained.  Hence 
arose  the  term  eudiometry^  signifying  to  measure  the  goodness. 
Cavendish,  on  the  other  hand,  as  the  result  of  a  large  number  of 
experiments  made  by  him,  came  to  the  conclusion  that  there  was 
no  difference  in  the  samples  of  air  that  he  experimented  upon. 

Since  the  time  of  Cavendish,  eudiometric  analysis  has  been 
brought  to  a  state  of  great  perfection  and  accuracy  by  Bunsen, 
Regnault,  Frankland,  and  others.  The  conclusion  to  be  drawn 
from  the  extended  researches  of  these  chemists  is,  that  although 
the  atmosphere  certainly  shows  a  remarkable  uniformity  of  com- 
position, there  do  exist  perceptible,  though  very  slight,  variations 

252 


The  Atmosphere 


253 


in  the  amount  of  oxygen  present  at  different  places  and  at  different 
times.     Samples  of  air  collected  from  all  parts  of  the  globe,  from 


mid  ocean,  irom  high  mountain  peak,  American  prairie,  and 
crowded  cities,  show  a  variation  in  the  proportion  of  oxygen  rang- 
ing from  20.99  t°  20.86.  Angus  Smith  has  shown  that  in  foggy 


254  Inorganic  Chemistry 

weather  the  oxygen  in  the  air  in  towns  sometimes  falls  as  low  as 
20.82.  Samples  of  air  taken  from  crowded  theatres  have  been 
found  to  contain  as  little  as  20.28,  while  in  many  mines  the  amount 
averages  as  low  as  20.26. 

The  mean  proportions  of  oxygen  and  nitrogen  in  the  atmosphere 
may  be  given  as — 

Oxygen  .....     20.96  parts  by  volume. 

Nitrogen*      ....     79-°4       »  » 

100.00 

The  composition  of  the  atmosphere  by  weight  was  determined 
by  Dumas  and  Boussingault  (1841).  In  their  method,  air  which  was 
freed  from  carbon  dioxide  and  moisture  was  slowly  drawn  through 
a  glass  tube  containing  a  known  weight  of  metallic  copper,  heated 
to  redness.  The  oxygen  combined  with  the  copper,  forming  copper 
oxide,  which  was  afterwards  weighed,  and  the  nitrogen  passed  into 
a  vacuous  flask,  and  was  also  weighed.  The  apparatus  as  em- 
ployed by  Dumas  is  seen  in  Fig.  53.  B  is  a  glass  flask  having  a 
capacity  of  10  to  15  litres,  which  was  exhausted  and  then  weighed. 
It  was  then  attached,  as  shown,  to  the  tube  T,  containing  a  known 
weight  of  metallic  copper,  and  which  was  also  exhausted.  The 
bulbs  L  contained  a  solution  of  potassium  hydroxide,  and  the  tubes 
/j  solid  potash,  for  the  removal  of  atmospheric  carbon  dioxide. 
The  bulbs  O  contained  strong  sulphuric  acid,  and  the  tubes  /  were 
filled  with  pumice  moistened  with  the  same  acid,  by  means  of 
which  the  moisture  was  withdrawn  from  the  air.  When  the  copper 
was  heated  and  the  cocks  partially  opened,  air,  free  from  carbon 
dioxide  and  moisture,  was  slowly  drawn  over  the  heated  metal, 
which  was  thereby  converted  into  the  oxide.  At  the  conclusion 
of  the  experiment  the  globe  and  the  tube  T  were  reweighed.  The 
nitrogen  remaining  in  tube  T  was  then  pumped  out  and  the  tube 
once  more  weighed.  The  difference  between  the  two  last  weigh- 
ings of  the  tube,  added  to  the  gain  in  weight  suffered  by  the  globe, 
gave  the  nitrogen  ;  while  the  difference  between  the  original  and 
final  weights  of  the  tube  gave  the  increase  of  weight  suffered  by 
the  copper,  that  is,  the  amount  of  oxygen.  The  result  of  numerous 
experiments  gave  the  mean  composition — 

Oxygen 23  parts  by  weight. 

Nitrogen  *          ....     77     „  „ 

100 
*  The  small  percentage  of  argon  present  is  here  included  with  the  nitrogen. 


The  Atmosphere 


255 


A  more  modern  method  for  estimating  the  amounts  of  oxygen 
and  nitrogen  in  the  air,  based  upon  the  same  principle,  namely,  the 
absorption  of  the  oxygen  by  heated  metallic  copper,  is  illustrated 
in  Fig.  54  (known  as  Jolly's  apparatus).  The  sample  of  air  to  be 
examined  is  allowed  to  enter  the  glass  globe  A  (whose  capacity  is 
about  100  c.c.,  and  which  has  been  previously  exhausted)  by  means 
of  the  three-way  cock  b.  (The  air  is  first  dried,  by  being  drawn 
through  tubes  filled  with  pumice  moistened  with  sulphuric  acid,  on 


FIG.  54. 

its  way  into  the  apparatus.)  The  bulb  is  then  surrounded  by  the 
metal  jacket  B,  which  is  filled  with  broken  ice,  and  when  the  tem- 
perature has  fallen  to  o°  the  bulb  is  put  into  communication  with 
the  tube  d  by  means  of  the  three-way  cock.  The  tube  g  is  then 
raised  or  lowered,  so  as  to  bring  the  mercury  in  d  to  a  fixed  point 
in  the  tube  at  ;#,  and  the  tension  of  the  enclosed  air  is  ascertained 
by  the  graduated  scale  behind  tube  g.  The  ice-jacket  is  then 
removed,  and  the  spiral  of  copper  wire  within  the  bulb  is  heated 
to  redness  by  the  passage  through  it  of  an  electric  current.  The 


356  Inorganic  Chemistry 

copper  combines  under  these  conditions  with  the  oxygen,  form- 
ing copper  oxide,  thereby  reducing  the  volume  of  the  contained 
gas.  The  globe  is  again  cooled,  and  the  tube  g  lowered  to  such 
a  position  that  when  communication  is  once  more  made  between 
the  globe  and  tube  d,  the  mercury  shall  stand  at  the  same  point  m. 
From  the  observed  tension  of  the  gas  before  and  after  the 
experiment,  the  volume  relations  of  the  two  constituents  can  be 
calculated.  Thus,  suppose  the  tension  of  the  enclosed  air  to  be 
720.25  mm.,  and  that  of  the  residual  nitrogen  569.28  mm.,  then  for 
I  volume  of  air  the  reduction  would  be  — 


Therefore  in  100  volumes  the  composition  would  be  — 
Nitrogen*  =  79.04 
Oxygen      =  20.96 


100.00 

Besides  oxygen  and  nitrogen,  the  air  contains  variable  quantities 
of  the  following  gases  :  aqueous  vapour,  carbon  dioxide,  argon, 
hydrogen,  ammonia,  ozone,  nitric  acid.  With  the  exception  of 
aqueous  vapour,  these  substances  are  present  only  in  relatively 
small  proportions,  and  with  some  of  them  the  amount  is  liable 
to  considerable  variation.  Especially  is  this  the  case  with  the 
aqueous  vapour,  as  the  amount  of  this  constituent  present  at  any 
time  is  largely  influenced  by  the  temperature.  The  average  com- 
position of  normal  air  may  be  taken  as  follows  : — 

Vols.  per  1000. 

Nitrogen 769.6500 

Oxygen         .         .         .         .         .  206.5940 

Aqueous  vapour   ....  14.0000 

Argon  t         .         .         «         .         .  9.3700 

Carbon  dioxide     ....  0.3360 

Hydrogen     .         .         .         .         .  0.0400 

Ammonia      .         .         .         .         .  0.0080 

Ozone   .         .         .         .         .         .  0.0015 

Nitric  acid    .....  0.0005 

IOOO.OOOO 

*  The  small  percentage  of  argon  present  is  here  included  with"  the  nitrogen, 
f  The  other  four  gases  of  the  argon  group  taken  together  come  to  about 
0.012  parts  per  1000  (see  page  270). 


The  Atmosphere  257 

Aqueous  Vapour. — For  any  given  temperature  there  is  a 
maximum  amount  of  aqueous  vapour  which  a  given  volume  of  air 
is  capable  of  taking  up  :  under  these  conditions  the  air  is  said  to 
be  saturated  with  moisture  at  the  particular  temperature.  Thus  I 
cubic  metre  of  air  is  saturated  with  moisture  at  the  various  tempe- 
ratures stated,  when  it  has  taken  up  the  following  weights  of 
water  : — • 


At  o°     .     4.871  grammes. 
„  10°     .     9.362 


At  20°    .     17.157  grammes. 
„  30°     .     30.095         „ 


When  air,  saturated  with  moisture  at  say  20°,  is  cooled  to  10°, 
the  excess  of  water  beyond  9.362  (the  maximum  for  10°)  is  deposited 
either  as  mist  or  rain.  The  temperature  at  which  air  thus  begins 
to  deposit  moisture  is  called  the  dew-point.  The  deposition  of 
moisture  from  the  air  caused  by  the  lowering  of  the  temperature 
is  a  matter  of  everyday  observation.  A  glass  vessel  containing 
iced  water  becomes  bedewed  with  moisture  upon  the  outside  as 
the  air  in  its  immediate  vicinity  is  cooled.  When  a  season  of 
severe  frost  is  suddenly  followed  by  a  warm  wind,  highly  charged 
with  aqueous  vapour,  it  is  not  unusual  to  see  condensed  moisture 
collecting  upon  and  streaming  down  the  cold  surface  of  walls. 
For  the  same  reason,  after  the  sun  has  set,  and  the  heat  from  the 
ground  has  radiated,  leaving  the  ground  colder  than  the  atmos- 
phere, the  temperature  of  the  air  is  lowered,  and  it  begins  to 
deposit  its  aqueous  vapour  in  the  form  of  dew. 

The  amount  of  aqueous  vapour  in  the  air,  or  the  humidity  of  the 
air,  is  estimated  by  meteorologists  by  means  of  an  instrument 
called  the  wet  and  dry  bulb  thermometer. 

Carbon  Dioxide. — The  proportion  of  this  gas  present  in  the 
air  is  also  liable  to  considerable  variation,  although  not  through 
such  a  wide  range  as  the  aqueous  vapour.  The  processes  of 
respiration,  combustion,  and  putrefaction  are  attended  by  the 
evolution  of  carbon  dioxide,  hence  the  amount  of  this  gas  present 
in  closed  inhabited  places  is  greater  than  that  in  the  open  air  ;  in 
badly  ventilated  and  crowded  rooms  the  proportion  sometimes 
rises  to  three  parts  in  1000  vols.  Frankland  has  found  that  at  high 
elevations  the  amount  of  carbon  dioxide  in  the  air  is  often,  although 
not  invariably,  considerably  above  the  normal. 

At  Chamounix  (3000  feet)  the  amount  of  carbon  dioxide  was  0.63  per  1000  vols. 
,,  Grands  Mulets  (11,000  feet)      ,,  ,,  „     i.n         ,,          „ 

„  Mont  Blanc  (15,732  feet)  ,,  „  „     0.61        „ 

R 


258  Inorganic  Chemistry 

This  fact  is  probably  due  to  the  absence,  in  these  high  regions, 
of  the  vegetation  which  is  one  of  the  chief  natural  causes  operating 
to  remove  atmospheric  carbonic  dioxide  (see  Oxygen,  page  188). 

The  amount  of  carbon  dioxide  is  slightly  higher  during  the 
night,  and  often  rises  considerably  during  foggy  weather.  Thorpe 
has  shown  that  near  the  surface  of  the  sea  the  amount  of  carbon 
dioxide  in  the  air  is  slightly  less,  being  on  an  average  0.300  volume 
per  looo. 

Ammonia  in  the  atmosphere  is  derived  from  the  decomposition 
of  nitrogenous  organic  matter.  Although  present  in  relatively 
very  small  quantities,  it  varies  in  amount  very  considerably.  From 
the  experiments  of  Angus  Smith,  1000  grammes  of  air  from  various 
sources  were  found  to  contain  the  following  amounts  of  ammonia  :  — 

London        .         .         .    0.05  gramme. 
Glasgow       .         .         .     0.06        „ 
Manchester.         .         .     o.io        „ 

The  proportion  of  ammonia  appears  to  be  higher  during  the 
night  than  in  the  daytime,  and  immediately  after  heavy  rain  the 
amount  is  perceptibly  diminished. 

Rain-water  always  contains  ammonia,  although  the  amount 
varies  greatly  with  changing  atmospheric  and  climatic  conditions. 
Lawes  and  Gilbert,  Angus  Smith,  and  others,  have  made  a  large 
number  of  estimations  of  the  amount  of  ammonia  in  rain-water  at 
various  places  and  seasons,  and  under  many  different  conditions. 

Nitric  Acid  is  produced  in  the  atmosphere  by  the  direct 
union  of  oxygen  and  nitrogen  whenever  a  lightning  flash  passes 
through  the  air  (see  Nitric  Acid).  Rain  which  falls  during  or 
immediately  after  a  thunderstorm  is  found  to  contain  nitrates  and 
nitrites. 

These  two  nitrogenous  compounds,  ammonia  and  nitric  acid, 
although  present  only  in  such  small  proportion  in  the  atmosphere, 
fulfil  a  most  important  function  in  the  economy  of  nature.  From 
the  experiments  of  Lawes  and  Gilbert,  and  others,  it  has  been  shown 
that  most  plants  are  unable  to  draw  upon  the  free  nitrogen  of  the 
atmosphere  for  the  supply  of  that  element  which  they  require  for 
the  development  of  their  structure  and  fruit.*  Although  they  are 
surrounded  by,  and  bathed  in  nitrogen,  they  cannot  assimilate  it. 
Plants  that  are  growing  in  unmanured  soil,  therefore,  derive  their 

*  Leguminous  plants,  such  as  clovers,  vetches,  beans,  peas,  which  develop 
root-nodules  or  tubercles,  are  exceptions. 


The  Atmosphere  259 

nitrogen  from  the  ammonia  and  nitric  acid  which  are  present  in 
the  air,  and  which  are  washed  into  the  ground  by  the  rain.  It  has 
been  found  that  a  plant  grown  under  such  experimental  conditions, 
as  to  exclude  the  possibility  of  its  obtaining  supplies  of  these  nitro- 
genous compounds,  will  yield  upon  analysis  exactly  the  same 
amount  of  nitrogen  as  was  originally  contained  in  the  seed  from 
which  it  grew. 

Ozone. — The  causes  which  operate  in  the  formation  of  this  sub- 
stance in  the  air  are  at  present  imperfectly  known  ;  it  is  supposed 
that ]  its  occurrence  is  related  to  the  development  of  electricity  in 
the  atmosphere.  On  account  of  the  powerful  oxidising  character 
of  ozone,  its  presence  can  never  be  detected  in  the  air  where  much 
organic  matter  of  an  oxidisable  nature  is  present,  as  is  the  case  in 
the  air  of  such  places  as  malarial  swamps,  dwelling-houses,  and 
large  towns. 

The  amount  of  ozone  in  pure  country  air  has  been  found  to  vary 
with  the  time  of  year,  reaching  a  maximum  in  the  spring-time,  and 
gradually  falling  towards  winter.  Thorpe  has  found  that  in  sea 
air  the  amount  of  ozone  is  practically  constant  jduring  all  seasons. 

The  usual  method  which  is  available  for  the  detection  and 
estimation  of  ozone  in  the  air  is  extremely  crude.  It  consists  in 
exposing  ozone  test  papers  (see  Ozone)  to  the  air  for  a  certain  time, 
and  comparing  the  colour  that  is  produced  with  a  standard  scale 
of  tints  ;  moreover,  other  substances  than  ozone,  which  may  be 
present  in  the  atmosphere,  will  also  liberate  iodine  from  potassium 
iodide,  and  these  are  therefore  measured  as  ozone.  Besides  the 
higher  oxides  of  nitrogen,  which,  as  we  have  seen,  are  formed  in  the 
atmosphere,  and  which  liberate  iodine  from  potassium  iodide,  it  has 
been  shown  that  peroxide  of  hydrogen  is  also  present.  The  state  of 
our  knowledge  at  present,  therefore,  respecting  the  exact  amount  of 
atmospheric  ozone  and  its  variation  is  far  from  satisfactory  ;  it  is, 
indeed,  quite  possible  that  many  of  the  effects  which  have  been  attri- 
buted to  ozone  are  in  reality  due  to  peroxide  of  hydrogen.  Thus  it 
has  been  shown  by  Schonbein  that  this  compound  is  formed  during 
the  evaporation  of  water,  and  this  statement  probably  derives  con- 
firmation from  the  fact  that  its  presence  may  be  detected  in  rain- 
water. The  salubrity  of  the  air  of  the  sea-shore,  where  large  areas 
of  wet  sand  and  stones  offer  the  most  perfect  conditions  for  the 
rapid  evaporation  of  water,  and  consequently,  for  the  formation  of 
peroxide  of  hydrogen,  may  therefore  be  attributable  as  much  to  the 
presence  of  this  substance  as  to  the  proverbial  ozone. 


260  Inorganic  Chemistry 

Hydrogen. — The  presence  of  this  gas  in  sensible  quantities  as 
a  constituent  of  normal  air  appears  to  have  escaped  notice  until 
quite  recently.  When  a  quantity  of  liquefied  air  is  subjected  to 
fractional  distillation,  the  first  and  most  volatile  portion  which 
collects  is  found  to  be  very  rich  in  hydrogen.* 

The  elaborate  researches  of  Gautiert  show  also  that  this  hydro- 
gen is  not  only  present  in  the  air  of  towns,  but  that  it  is  a  con- 
stituent of  country  air,  air  from  high  mountain  regions,  and  of  sea 
air.  The  chief  sources  of  this  hydrogen  are  indicated  on  page 
171. 

The  various  gases  of  which  the  air  is  composed  are  not  com- 
bined, but  are  merely  mingled  together.  The  remarkable  con- 
stancy of  its  composition,  as  regards  the  oxygen  and  nitrogen,  led 
chemists  at  one  time  to  suppose  that  these  gases  were  in  chemical 
union  with  each  other  in  the  atmosphere ;  but  a  number  of  facts 
which  have  since  been  learnt  respecting  these  gases  prove  with- 
out doubt  that  this  is  not  the  case,  and  that  the  air  is  simply  a 
mechanical  mixture.  This  evidence  may  be  briefly  summed  up 
as  follows  : — 

(i.)  When  oxyg<?n  and  nitrogen  are  mixed  together  in  the 
proportion  in  which  they  occur  in  air,  the  resulting  mixture 
behaves  in  all  respects  like  ordinary  air,  and  the  mixing  of  the 
gases  is  not  attended  by  any  volumetric  or  thermal  disturbance, 
such  as  would  be  expected  to  accompany  the  chemical  union  of 
two  elements. 

(2.)  The  degree  to  which  air  is  capable  of  refracting  light  is 
found  to  be  the  mean  of  the  refractive  power  of  oxygen  and 
nitrogen.  Were  these  gases  chemically  combined,  the  compound 
should  behave  in  this  respect  as  other  compound  gases,  where  it 
is  found  that  the  refractive  index  is  always  either  greater  or  less 
than  the  mean  of  that  of  the  constituents. 

(3.)  According  to  a  fundamental  law  of  chemical  science,  the 
composition  of  a  chemical  compound  is  constant.  Such  a  thing  as 
variability  in  the  composition  of  a  compound  is  unknown.  The 
proportion  of  oxygen  and  nitrogen,  as  we  have  seen,  does  vary  in 
the  air,  although  through  only  small  limits,  hence  they  cannot  be 
united  to  form  a  compound. 

(4.)  The  proportion  by  weight  in  which  oxygen  and  nitrogen  are 

*  Dewar,  Nature,  December  20,  1900. 

f  Gautier,  Annales  de  Chimie  et  de  Physique,  January  1901. 


The  Atmosphere  261 

present  in  air  bears  no  simple  relation  to  the  atomic  weights  of 
these  elements. 

(5.)  When  air  is  dissolved  in  water,  the  oxygen  and  nitrogen 
dissolve  as  from  a  simple  mixture  of  these  gases,  in  accordance  to 
the  law  of  partial  pressures  (see  page  147). 

(6.)  The  oxygen  and  nitrogen  can  be  partially  separated,  by 
taking  advantage  of  the  different  rates  of  diffusion  of  these  two 
gases  (see  Diffusion  of  Gases,  page  83). 

The  various  gases  of  the  atmosphere  are  maintained  in  a  state 
of  uniform  admixture,  in  spite  of  their  widely  different  densities, 
by  the  operation  of  two  causes  :  first,  air  currents,  which  effect  the 
rapid  removal  of  large  masses  of  air  from  place  to  place  ;  and, 
second,  their  own  molecular  movements,  which  bring  about  the 
phenomena  of  gaseous  diffusion. 

Suspended  Impurities  in  the  Atmosphere.— Besides  the 
gaseous  constituents  of  the  air,  there  is  always  present  a  certain 
quantity  of  suspended  matter,  both  liquid  and  solid.  The  exist- 
ence of  this  suspended  matter  in  the  air  can  be  rendered  evident 
from  the  fact  that  these  minute  particles  are  capable  of  reflecting 
light ;  if,  therefore,  a  strong  beam  of  light  be  passed  through  a 
darkened  room,  the  track  of  the  beam  is  distinctly  visible,  on 
account  of  its  being  reflected  from  innumerable  particles  floating 
about  in  the  air,  many  of  them  appearing  quite  large.  Pasteur  has 
shown  that  this  suspended  matter  can  be  removed  by  filtration 
through  cotton  wool.*  Tyndall  also  has  shown  that  in  undis- 
turbed air  the  suspended  matter  settles  in  the  course  of  a  few 
hours,  leaving  the  air  almost  entirely  free  from  this  impurity. 
For  this  purpose  the  floor  of  a  large  oblong  glass  box  was' 
smeared  over  with  glycerine.  The  box,  after  being  hermetically 
closed,  was  then  allowed  to  stand  for  twenty-four  hours,  during 
which  time  the  suspended  matter  subsided  and  adhered  to  the 
glycerine.  When  a  beam  of  light  is  allowed  to  pass  through  air 
that  has  been  thus  freed  from  suspended  matter,  there  being 
nothing  present  to  reflect  the  light,  the  beam  cannot  be  seen  ; 
its  track  will  be  evident  in  the  air  of  the  room  as  it  enters  and 
leaves  the  box,  but  within  the  box  it  will  be  invisible  (as  repre- 
sented in  Fig.  55).  To  air  in  which  a  beam  of  light  is  in  this  way 
invisible,  Tyndall  has  applied  the  term  "  optically  pure." 

The  suspended  matters  are  partly  mineral  and  partly  organic. 
Of  the  mineral  matters,  sodium  chloride  and  certain  sulphates 

*  See  Experiments  334  to  341,  "  Chemical  Lecture  Experiments,"  new  ecu 


262  Inorganic  Chemistry 

are  present  in  greatest  quantity.  These  are  thrown  into  the  air 
in  the  sea-spray,  and  as  the  small  globules  of  water  evaporate 
they  leave  minute  residual  particles  of  saline  matter,  which, 
being  driven  by  the  wind,  remain  floating  in  the  atmosphere.  It 
is  only  very  rarely,  even  at  far  inland  places  in  Europe,  that  spec- 
troscopic  examination  fails  to  detect  the  presence  of  sodium  com- 
pounds in  the  air.  In  the  air  of  islands,  such  as  England,  it  is 
never  absent.  Sulphates  are  also  produced  by  the  oxidation  and 
combustion  of  sulphuretted  compounds  ;  the  amount  of  these,  there- 
fore, is  greatly  increased  in  the  neighbourhood  of  towns. 

The  organic  suspended  matter  of  the  air  has  of  late  years  been 
made  the  subject  of  extended  research.  Pasteur  has  shown  that 
amongst  these  organic  substances  are  the  germs  and  organisms 
which  produce  fermentation,  putrefaction,  and  disease.  Putrescible 


FIG.  55. 

substances,  such  as  milk,  urine,  flesh,  &c.,  if  themselves  carefully 
freed  from  all  such  germs,  may  be  preserved  unchanged,  for  ap- 
parently any  length  of  time,  in  air  that  has  been  deprived  of  all 
suspended  matter.  It  is  highly  probable  that  the  salubrity  or 
otherwise  of  different  places  is  associated  with  the  nature  and 
amount  of  the  organic  matter  in  the  air,  and  it  is  certain  that 
these  organisms  play  a  most  important  part  in  relation  to  the  life 
and  health  of  man.  The  feelings  of  lassitude  and  headache,  which 
result  from  the  prolonged  breathing  of  the  air  of  rooms  containing 
many  people,  are  brought  about  more  by  the  poisonous  effects  of 
the  organic  emanations  evolved  during  respiration  than  by  any 
diminution  in  the  supply  of  oxygen,  or  increase  in  the  proportion 
of  carbon  dioxide  in  the  air.  The  well-known  and  unpleasant 
smell  that  is  perceived  on  first  entering  a  crowded  room  is  also 
due  to  the  same  cause,  and  it  has  been  shown  that  the  moisture 
which  condenses  from  such  an  atmosphere  upon  a  cold  object,  if 


The  Atmosphere  263 

preserved  for  a  short  time,  rapidly  becomes  putrescent,  owing  to 
the  decomposition  of  this  organic  matter. 

The  presence  of  suspended  matter  in  the  air  appears  to  exert  a 
remarkable  influence  upon  the  formation  and  character  of  fogs. 
Aitkin  has  shown  that  those  conditions  which  result  in  the  forma- 
tion of  a  fog  in  ordinary  air  are  incapable  of  producing  that  effect 
in  air  that  has  been  freed  from  suspended  matter.  It  would  appear 
that  the  suspended  particles  act  as  innumerable  points,  or  nuclei, 
which  facilitate  the  deposition  of  moisture,  much  in  the  same  way 
as  the  crystallisation  of  a  salt,  from  its  solution,  is  known  to  start 
from  any  minute  particles  of  foreign  matter  that  may  be  floating  in 
the  liquid. 

The  height  to  which  the  atmosphere  extends  has  been  variously 
estimated.  From  observation  of  the  flight  of  meteorites,  it  ap- 
pears that  even  at  a  height  of  seventy  to  seventy-five  miles  the 
air  still  has  a  sensible  degree  of  density.  The  air  being  elastic, 
and  subject  to  the  law  of  gravitation,  its  density,  which  is  greatest 
at  the  earth's  surface,  rapidly  diminishes  as  the  altitude  increases  ; 
thus,  at  about  three  and  a  half  miles  the  density  is  only  one-half, 
and  at  seven  miles  one-third,  of  that  which  obtains  at  the  sea-level. 
From  a  consideration  of  the  physical  properties  of  gases,  there  is 
every  reason  to  believe  that  in  an  extremely  attenuated  condition 
the  atmosphere  extends  far  into  space,  and  it  has  been  calculated 
that  the  pressure  exerted  by  our  atmosphere  upon  the  surface  of 
the  moon  is  equal  to  about  I  mm.  of  mercury. 

The  density  of  the  atmosphere  varies  at  different  points  of  the 
earth's  surface,  and  at  the  same  point  at  different  times.  The 
pressure  exerted  by  the  atmosphere  is  measured  by  the  height 
of  a  column  of  mercury  which  it  is  capable  of  supporting,  the 
instrument  employed  for  the  purpose  being  called  the  barometer. 
At  the  sea-level  in  the  latitude  of  London,  the  average  weight  of 
the  atmosphere  is  equal  to  that  of  a  column  of  mercury  760  mm. 
at  o°,  and  this  is  taken  as  the  standard  pressure  of  the  atmos- 
phere. 

THE  ARGON  GROUP  OF  ATMOSPHERIC  GASES 

History. — More  than  a  hundred  years  ago  Cavendish  observed 
that  when  a  mixture  of  nitrogen  (phlogisticated  air)  and  oxygen 
(dephlogisticated  air)  was  confined  in  a  glass  tube  over  mercury 
along  with  a  solution  of  caustic  potash,  and  the  gases  exposed  to 


264  Inorganic  Chemistry 

the  continued  action  of  electric  sparks,  there  was  a  small  residue 
of  gas  (amounting  to  about  T^th  of  the  volume  of  the  nitrogen) 
which  was  not  absorbed,  and  he  raised  the  question  as  to  whether 
the  "  phlogisticated  air"  of  our  atmosphere  is  entirely  of  one  kind.* 

This  observation  and  speculation  of  Cavendish's  remained  buried 
until  1894,  when  Lord  Rayleigh  and  Professor  Ramsay  announced 
to  the  world  the  discovery  of  a  new  gaseous  constituent  of  the 
atmosphere. 

In  making  exact  determinations  of  the  densities  of  gases,  Lord 
Rayleigh  found  that  nitrogen  obtained  from  atmospheric  sources 
always  gave  a  slightly  higher  number  than  that  obtained  for 
nitrogen  which  was  prepared  from  chemical  compounds.  On 
careful  investigation,  in  conjunction  with  Professor  Ramsay,  it  was 
found  that  this  higher  density  of  "  atmospheric  nitrogen  "  was  due 
to  the  presence  in  the  air  of  a  hitherto  unknown  gas,  which  they 
succeeded  in  isolating,  and  to  which  they  gave  the  name  Argon 

(1894). 

In  the  following  year,  in  searching  for  probable  sources  of 
argon,  Ramsay  was  led  to  examine  the  gas  which  was  known  to 
be  occluded  in  certain  minerals,  notably  in  the  rare  minerals 
cleveite  and  broggerite.  This  gas,  which  had  hitherto  been  regarded 
as  nitrogen,  was  found  to  give  a  spectrum  the  most  characteristic 
line  of  which  was  a  remarkably  brilliant  one  in  the  yellow.  The 
position  of  this  yellow  line  proved  to  be  coincident  with  the 
line  D3  of  the  solar  spectrum,  which  is  the  characteristic  line  of  a 
hitherto  unknown  solar  element  first  observed  by  M.  Janssen  of 
Paris  in  1868,  the  spectrum  of  which  was  studied  by  Frankland 
and  Lockyer,  who  applied  the  name  "helium"  ("the  sun")  to  the 
element.  Subsequently  Ramsay  has  shown  that  helium  is  present 
in  the  atmosphere,  although  in  much  smaller  quantities  than  argon. 

In  the  year  1898  the  discoverer  of  argon  announced  to  the  Royal 
Society  the  discovery  of  two  other  gases  which  were  associated 
with  argon,  to  which  he  gave  the  names  neon  ("the  new  one") 
and  krypton  ("the  hidden  one"),  and  subsequently  he  discovered 
still  another  and  denser  gas,  which  has  been  called  xenon  ("  the 
stranger").  The  five  gases,  therefore,  belonging  to  this  group,  in 
the  order  of  their  densities,  are  :  helium,  2  ;  neon,  10;  argon,  19.95  ; 
krypton,  41.50;  xenon,  65.35. 

Not  only  are  these  five  new  gases  elementary  substances,  but 
they  possess  many  properties  in  common.  They  all  are  extremely 

*  "Experiments  on  Air,"  Phil.  Trans.,  75,  372,  1785. 


Argon  265 

inert  elements,  having  apparently  no  chemical  activities  whatever. 
No  compounds  are  known  in  which  any  one  of  them  exists  as  a 
chemical  constituent,  and  they  have  resisted  all  attempts  to  cause 
them  to  enter  into  chemical  combination  with  any  other  element. 
Hence  there  is  at  present  no  chemistry  of  these  strange  substances. 
This  being  the  case,  the  only  light  which  can  be  thrown  upon  the 
complexity  of  the  molecules  of  these  elements  is  by  the  determina- 
tion of  the  ratio  of  their  specific  heats  at  constant  pressure  and  at 
constant  volume,  deduced  from  determinations  of  the  wave-length 
of  sound.  This  ratio  is  found  to  be  1.66,  which  is  the  same  as  that 
obtaining  in  the  case  of  mercury  vapour,  the  only  other  monatomic 
gas  in  which  this  ratio  has  been  determined  ;  whereas  with  diatomic 
gases,  such  as  hydrogen,  nitrogen,  and  oxygen,  the  ratio  is  1.4. 

AEGON. 

Symbol,  A.     Density,  19.95.     Atomic  weight,  39.9. 

Occurrence.—  Argon  is  present  in  the  atmosphere,  where  it 
exists  to  the  extent  of  0.937  per  cent,  or  rather  more  than  i  per 
cent,  of  the  "atmospheric  nitrogen"  is  argon.  It  is  also  found  in 
the  occluded  gases  of  certain  specimens  of  meteoric  iron,  and  in 
minute  quantities  in  almost  all  natural  waters,  derived  doubtless 
by  solution  from  the  atmosphere.  Argon  has  not  been  met  with 
in  chemical  combination  with  other  elements,  and  no  compounds 
containing  this  element  are  known. 

Modes  Of  Preparation.— ( i.)  Argon  may  be  obtained  from  the 
atmosphere  by  sparking  a  mixture  of  air  and  oxygen.  The  nitrogen 
combines  with  the  oxygen,  and  the  oxidised  product  is  absorbed 
by  potash.  When  no  further  contraction,  of  volume  is  obtained, 
the  excess  of  oxygen  is  removed  by  alkaline  pyrogallate,  and  the 
residual  gas  is  the  argon.  Unless  a  high-tension  alternating 
electric  discharge  is  employed  the  process  is  extremely  slow. 

(2.)  Argon  may  also  be  separated  from  the  other  atmospheric 
gases  by  first  withdrawing  the  oxygen  by  means  of  red-hot  copper, 
and  after  removing  the  carbon  dioxide  and  aqueous  vapour,  passing 
the  remaining  gas  over  strongly  heated  magnesium  turnings.  The 
magnesium  combines  with  the  nitrogen  (p.  232)  and  leaves  the 
argon.  In  order  to  effect  the  complete  absorption  of  every  trace 
of  nitrogen,  the  gas  is  passed  backwards  and  forwards  over  the 
heated  magnesium  for  many  hours. 

Recently  it  has  been  found  that  the  metal  calcium  is  a  more 


266  Inorganic  Chemistry 

efficient  agent  for  the  absorption  of  nitrogen.  If,  therefore,  the 
"  atmospheric  nitrogen  "  be  passed  over  a  heated  mixture  of  mag- 
nesium filings  and  pure  dry  lime  the  magnesium  and  lime  interact, 
forming  magnesia  and  calcium,  which  latter  absorbs  the  nitrogen 
very  rapidly  and  at  a  lower  temperature  than  that  required  by 
metallic  magnesium. 

The  purification  of  the  argon  thus  obtained,  and  its  complete 
separation  from  the  other  gases  with  which  it  is  associated,  was  a 
problem  which  was  only  solved  as  a  result  of  the  later  achievements 
by  Dewar  of  obtaining  liquid  hydrogen  in  quantity.  By  means  of 
the  intense  cold  obtainable  by  liquid  hydrogen,  comparatively 
large  quantities  of  argon  were  liquefied,  and  the  liquid  so  obtained 
was  then  submitted  to  a  process  of  fractional  distillation.  The 
liquid  gases  having  the  lowest  boiling-points,  namely  helium  and 
neon,  are  the  first  to  evaporate,  and  by  careful  adjustment  of  the 
temperature  of  the  refrigerating  bath  the  denser  gases,  krypton 
and  xenon,  may  be  maintained  even  in  the  solidified  state,  while  the 
whole  of  the  argon  in  a  state  of  practical  purity  can  be  distilled  away. 

Properties.— Argon  is  remarkable  for  its  extraordinary  inert- 
ness, a  property  which  is  indicated  by  its  name,  argon  signifying 
"  inactive."  As  already  mentioned,  it  has  hitherto  resisted  all 
attempts  to  cause  it  to  unite  chemically  with  any  other  element. 
The  density  of  the  gas  is  19.95,  and  therefore  its  molecular  weight 
is  39.9  ;  and  since  argon  is  a  monatomic  element,  its  atomic 
weight  and  its  molecular  weight  are  the  same. 

Argon  is  about  two  and  a  half  times  as  soluble  in  water  as  nitro- 
gen, 100  volumes  of  water  at  15°  dissolving  4.1  volumes  of  argon. 
Owing  to  this  superior  solubility,  the  gases  which  are  expelled 
from  rain-water  by  boiling  are  slightly  richer  in  argon  than  the 
original  air  before  solutibn.  The  critical  temperature  of  argon  is 
— 117.4°,  at  which  temperature  the  gas  is  liquefied  by  a  pressure  of 
about  fifty-three  atmospheres  (or  40.20  metres  of  mercury).  Liquid 
argon  has  a  specific  gravity  of  1.212,  and  boils  at  -  186.1°.  The 
boiling-point  of  argon,  therefore,  lies  between  those  of  the  two  chief 
constituents  of  the  atmosphere,  namely,  oxygen  and  nitrogen,  while 
its  critical  temperature  is  slightly  above  that  of  oxygen,  as  may  be 
seen  by  the  following  comparison  : — 

Boiling-point.          Critical  Temperature. 

Oxygen   ....     -182.5°  -118.8° 

Argon      ....     -186.1°  -117-4° 

Nitrogen         ...      - 195.5°  - 149° 


Helium  267 

A  very  slight  reduction  of  temperature  below  its  boiling-point  is 
sufficient  to  freeze  argon  to  a  white  solid,  the  melting-point  of 
which  is  —  187.9°  ;  that  is,  less  than  two  degrees  below  the  boiling- 
point.  The  spectrum  of  argon  is  very  complex.  The  most  cha- 
racteristic lines  are  two  in  the  red  (less  refrangible  than  the  red 
lines  of  either  hydrogen  or  lithium),  a  bright  yellow  line  (more 
refrangible  than  the  sodium  line),  a  group  of  bright  green  lines, 
and  another  group  of  strong  lines  in  the  violet.  The  general  cha- 
racter of  the  spectrum  depends  upon  the  nature  of  the  electric 
discharge  employed.  With  an  intermittent  discharge  the  lines  in 
the  red  and  pale  green  are  the  most  prominent,  while  with  a 
Leyden-jar  discharge  the  red  and  light  green  lines  almost  entirely 
disappear,  giving  place  to  lines  in  the  dark  green,  blue,  and  violet. 

HELIUM. 

Symbol,  He.     Density.  2.     Atomic  weight,  4. 

Oeeurrenee. — The  existence  of  this  element  in  the  universe 
may  be  said  to  have  been  first  discovered  by  Janssen,  who  during 
a  solar  eclipse  in  1868  observed  a  certain  line  in  the  yellow  of  the 
spectrum  of  the  sun's  chromosphere  which  was  not  coincident  with 
that  of  any  known  terrestrial  element.  This  unknown  element 
was  afterwards  named  helium  by  Frankland  and  Lockyer. 

Terrestrial  helium  was  discovered  by  Ramsay  in  1895,  in  the  gas 
which  is  contained  in  certain  rare  minerals,  and  which  is  evolved 
from  them  either  when  they  are  heated  or  when  they  are  treated 
with  dilute  sulphuric  acid.  Chief  among  these  minerals  are 
cleveite,  broggerite,  and  uraninite,  all  of  them  minerals  containing 
the  metal  uranium.* 

Helium  is  present  in  minute  quantities  in  the  atmosphere, 
namely,  to  the  extent  of  about  I  or  2  volumes  in  1,000,000  volumes 
of  air,  as  estimated  by  its  discoverer. 

Helium  also  occurs  in  certain  natural  waters,  notably  in  the 
water  from  the  Bath  springs,  which  has  been  found  to  contain 
argon  mixed  with  about  8  per  cent,  of  its  volume  of  helium. 

Method  Of  Preparation.— Helium  is  isolated  from  the  atmos- 
phere by  a  method  consisting  firstly  of  what  may  be  described  as 
fractional  liquefaction,  followed  by  fractional  evaporation  or  distilla- 
tion. When  air  is  liquefied  by  the  so-called  self-cooling  or  recupe- 

*  The  amount  of  helium  contained  in  i  gramme  of  cleveite  is  about  3.2  c.c. 
(Ramsay),  only  about  one  half  of  which  is  given  off  by  heat  alone. 


268  Inorganic  Chemistry 

rative  method  in  any  of  the  modern  air-liquefiers  based  upon 
Linde's  original  apparatus  (see  page  77),  those  portions  of  the  air 
which  escape  liquefaction  and  pass  out  of  the  apparatus  will 
obviously  contain  most  of  the  constituents  having  the  lowest  boil- 
ing-points. Therefore,  by  collecting  the  gas  which  escapes  from 
the  air-liquefier  under  these  circumstances,  and  compressing  it  into 
a  vessel  cooled  by  liquid  air,  a  liquid  is  obtained  which  contains 
most  of  the  more  volatile  constituents  (namely,  the  helium  and 
neon),  with,  of  course,  argon  and  some  nitrogen.  Thus,  by  this 
process  of  fractional  liquefaction  liquid  air  is  divided  into  two 
fractions,  one  containing  practically  all  the  denser  and  least  vola- 
tile constituents,  namely,  the  krypton  and  xenon,  the  other  con- 
taining the  helium  and  neon. 

The  separation  of  the  gases  contained  in  the  more  volatile 
fraction  is  accomplished  by  fractional  distillation  or  evaporation. 
At  the  low  temperature  obtainable  by  means  of  liquid  hydrogen 
both  argon  and  neon  exert  no  vapour-pressure,  being  reduced  to 
the  state  of  non-volatile  solids,  and  the  helium  in  a  state  of  purity 
can  be  pumped  away  from  the  mixture. 

Properties. — Next  to  hydrogen,  helium  is  the  lightest  known 
gas,  its  density  being  2.  Like  all  the  other  gases  of  this  group 
its  molecules  are  monatomic,  its  atomic  and  molecular  weight 
therefore  is  4.  Helium  is  much  less  soluble  in  water  than 
argon.  The  solubility  of  this  gas  in  water  forms  an  exception  to 
the  usual  behaviour  of  gases,  for  it  has  been  found  that  while  its 
solubility  diminishes  with  rise  of  temperature  up  to  about  25°, 
between  25°  and  50°  the  solubility  slightly  increases,  as  is  shown 
by  the  following  table.* 

100  volumes  of  water  at  760  mm.  dissolve — 

o°  10°  20°  30°  40°  50* 

0.01500        0.01442        0.01386        0.01382        0.01387        0.01404  vols. 

When  induction  sparks  are  passed  through  rarefied  helium,  the  gas 
emits  a  brilliant  yellow  light  with  a  tinge  of  apricot  colour.  When 
viewed  through  the  spectroscope  the  most  prominent  and  charac- 
teristic line  is  the  intense  yellow  line  D3,  which  is  accompanied  by 
one  bright  red  line,  two  in  the  green,  and  two  in  the  blue.  On 
reducing  the  pressure  in  the  tube,  the  yellow  light  due  to  line  D3 
gradually  changes  to  a  green,  owing  to  the  light  from  one  of  the 
green  lines  becoming  greatly  intensified. 

*  Estreicher,  "  Z.  Phys.  Ch.,"  31,  176. 


Neon  269 

Helium  is  the  last  of  the  known  gases  to  be  liquefied,  having 
resisted  all  attempts  upon  it  until  July  1908,  when  its  liquefac- 
tion was  successfully  accomplished  by  Prof.  Omnes.  200  litres  of 
the  purest  helium  after  circulating  for  some  hours  through  the 
"liquefier"  cooled  by  liquid  hydrogen  boiling  under  reduced  pres- 
sure, yielded  as  much  as  60  c.c.  of  liquid  helium.  Liquid  helium 
boils  about  -268.7°  or  4.3°  absolute,  its  critical  temperature  being 
about  5°  absolute. 

Lijte  all  the  other  gases  of  this  group  helium  is  chemically  inactive. 

NEON. 

Symbol,  Ne.     Density  =  10.     Atomic  weight  =  20. 

History. — From  analogy  with  other  natural  families  of  elements 
and  the  numerical  relations  of  the  atomic  weights  of  the  different 
members,  the  discoverer  of  argon  and  helium  was  led  to  believe 
that  another  element  should  exist  having  an  atomic  weight  between 
those  of  these  two  elements,  and  about  sixteen  units  higher  than 
that  of  helium.  The  long  and  careful  search  for  this  unknown 
element  was  at  last  rewarded  by  the  discovery  of  neon,  whose 
atomic  weight  was  found  to  be  20,  or  exactly  sixteen  units  above 
that  of  helium. 

Although  only  present  in  minute  quantities  in  the  atmosphere, 
the  discoverer  estimates  the  amount  as  about  ten  times  that  of 
helium  ;  that  is  to  say,  I  or  2  parts  of  neon  in  100,000  parts  of  air. 

Neon  is  more  readily  liquefied  than  helium,  but  no  exact  deter- 
mination of  its  boiling-point  has  yet  been  made. 

The  colour  emitted  by  this  gas  when  induction  sparks  are  passed 
through  it  is  a  brilliant  orange-pink.  Its  spectrum  is  characterised 
by  a  bright  yellow  line,  D6,  and  a  great  cluster  of  lines  in  the  more 
orange  part  of  the  red  end.  It  also  exhibits  other  fainter  lines 
throughout  the  spectrum. 

KRYPTON  AND  XENON. 

Krypton,  symbol  Kr.     Density  =  41. 50.      Atomic  weight  =  83.0. 
Xenon  ,,       X.  ,,      =65.35.  »  »       =I3°-7- 

These  two  denser  gases  are  obtained  by  the  fractional  distillation 
of  the  heavier  portion  of  liquid  air  obtained  in  the  air-liquefier 
(see  Helium).  Large  quantities  of  this  liquid,  amounting  to  30 
litres,  were  carefully  evaporated,  and  the  residual  portion,  after 
being  entirely  freed  from  nitrogen  and  oxygen,  was  again  liquefied 
by  means  of  liquid  air.  The  constituents  of  this  liquid  were  then 
separated  by  fractionation.  As  soon  as  most  of  the  argon  was 


270  Inorganic  Chemistry 

removed,  the  residue  consisting  of  krypton  and  xenon  was  easily 
solidified.  Under  these  circumstances  it  was  found  that  the  krypton 
could  be  withdrawn  by  pumping,  for  at  the  temperature  of  liquid 
air  solid  krypton  is  appreciably  volatile,  while  the  solidified  xenon 
is  practically  non-volatile.  In  the  estimation  of  the  discoverer,  air 
contains  only  about  one  part  of  krypton  in  one  million  parts  ; 
while  of  xenon  the  proportion  is  about  one  part  in  twenty  millions. 
The  boiling  and  melting  points  of  krypton  are  —  151.7°  and  —  169°  re- 
spectively, while  those  of  xenon  were  found  to  be  —  109.1°  and  —  140°. 

The  light  emitted  by  krypton,  under  the  influence  of  the  induc- 
tion spark,  is  a  yellowish-green  colour,  while  that  given  by  xenon 
under  the  same  circumstances  is  more  of  a  sky-blue. 

The  most  prominent  lines  in  the  spectrum  of  krypton  are  two 
very  near  together  in  the  red,  one  bright  yellow  line  and  one  strong 
green  line  ;  besides  which  there  are  a  few  in  the  blue  and  violet. 
The  brilliant  green  line  (wave  length,  5570.5)  has  attracted  special 
notice,  as  it  is  considered  highly  probable  that  this  line  may  prdve 
to  be  coincident  with  the  chief  line  in  the  aurora  spectrum. 

The  spectrum  of  xenon,  like  that  of  argon,  is  markedly  different 
as  the  electric  discharge  is  modified.  With  the  intermittent  dis- 
charge the  prominent  lines  are  four  in  the  red  end,  and  a  number 
of  strong  lines  in  the  blue  and  greenish-blue.  With  the  "jar" 
discharge  the  red  and  biue  lines  become  very  reduced  or  alto- 
gether disappear,  and  their  place  is  taken  by  a  number  of  lines  in 
the  bright  green. 

The  relative  proportion  in  which  these  gases  are  believed  to  be 
present  in  the  atmosphere  has  been  estimated  provisionally  by  the 
discoverer  as  follows  : — 

Helium,    I  part  per      250,000  of  air. 

Neon,        i  to  2  parts    „         100,000       „ 
Argon,      0.937  part      „  100       „ 

Krypton,  I  „        „     1,000,000       „ 

Xenon,     I  „        „  20,000,000       „ 

Or,  expressed  in  parts  per  1000,  to  compare  more  readily  with 
the  figures  in  the  table  on  page  256  : — 

Argon 9.37 

Neon o.oi 

Helium        ....  0.004 

Krypton       .  o.ooi 

Xenon .....  0.00005 


The  Inert  Gases 


271 


The  following  table  gives  the  latest  physical  constants  for  the 
members  of  this  strange  group  of  new  elements. 


1  ><..-n- 
sity. 

Atomic 
Weight. 

Boiling- 
Point. 

Melting- 
Point. 

Critical 
Temp. 

Criti-ul 
Pressure. 

Helium 

2 

4 

-268.7° 

-268° 

Neon 

10 

20 

..: 

Argon   . 

19-95 

39-9 

-186.1° 

-187.9° 

-117.4° 

52.9  Ats. 

Krypton 

*4i-S 

83.0 

-151-7° 

-169° 

-    62.5° 

54-2    ,, 

Xenon  . 

*65.35 

130.7 

—  109.  i° 

-I40° 

+    I4-75° 

57-2    ,, 

The  members  of  this  group,  while  exhibiting  many  close  re- 
semblances, such  as  the  monatomic  nature  of  their  molecules, 
their  remarkable  inertness,  &c.,  show  also  that  gradation  of  pro- 
perties which  is  met  with  in  other  natural  groups  of  elements. 
This  appears  by  the  results  tabulated  above,  as  well  as  by  such 
other  properties  as  the  refractive  indices,  atomic  volumes,  &c. 


Moore  and  Ramsay,  Jour.  Chem.  Soc. ,  December  1908. 


CHAPTER  VII 
COMPOUNDS  OF  NITROGEN  AND  HYDROGEN 

FOUR  compounds  of  nitrogen    with   hydrogen   have    been    pre- 
pared, namely : — 

Ammonia        .....  NH3. 

Hydrazine N2H4  or  (NH2)2. 

Hydrazoic  acid       ....  N3H  or  HN3. 

Ammonium  hydrazoate  .         .  N4H4  or  NH4N3. 

AMMONIA. 

Formula,  NH3.     Molecular  weight =17. 04.     Density=8.52. 

History. — Ammonium  salts,  and  also  the  aqueous  solution  of 
ammonia,  were  known  to  the  alchemists.  It  was  termed  by 
Glauber,  spiritus  volatilis  salts  armoniaci,  being  obtained  by  the 
action  of  an  alkali  upon  sal-armoniacum.  Subsequently  the  name 
spirits  of  hartshorn  was  applied  to  the  ammoniacal  liquid  obtained 
by  the  destructive  distillation  of  such  refuse  as  hoofs  and  horns  of 
animals.  The  actual  discovery  of  gaseous  ammonia  was  made  by 
Priestley  (1774),  when  he  collected  the  gas,  evolved  by  the  action 
of  lime  upon  sal-ammoniac,  by  means  of  his  mercurial  pneumatic 
trough.  Priestley  named  the  gas  alkaline  air. 

Occurrence.— In  combination  as  ammonium  carbonate  it  is  pre- 
sent in  small  quantities  in  the  air,  derived  by  the  decay  of  nitro- 
genous animal  and  vegetable  matter.  As  nitrate  and  nitrite  it  is 
found  in  rain-water.  It  is  evolved,  along  with  boric  acid,  from  the 
fumaroles  of  Tuscany  (see  Boric  Acid),  and  is  found  as  ammonium 
chloride  and  sulphate  in  the  vicinity  of  active  volcanoes. 

Modes  of  Formation. — (i.)  Ammonia  can  be  synthetically  pro- 
duced by  submitting  a  mixture  of  nitrogen  and  hydrogen  to  the 
influence  of  the  silent  electric  discharge  (Donkin).  The  amount 
of  ammonia  so  obtained,  however,  is  extremely  small,  and  can  best 
be  shown  by  passing  the  gases,  as  they  issue  from  the  "ozone 
tube,"  through  a  cylinder  containing  a  small  quantity  of  Nessler's 

272 


Ammonia  273 

solution.*  In  a  short  time  the  solution  will  begin  to  show  a 
yellowish-brown  colour,  indicating  the  presence  of  traces  of 
ammonia. 

(2.)  Ammonia  may  be  prepared  by  gently  heating  any  of  the 
ammonium  salts,  with  either  of  the  caustic  alkalies,  potash  or 
soda,  or  with  slaked  lime.  The  salt  most  commonly  employed  is 
the  chloride.  When  this  is  mixed  with  an  excess  of  slaked  lime, 
and  the  mixture  gently  heated  in  a  flask,  ammonia  is  evolved,  and 
calcium  chloride  and  water  are  formed— 


4Cl  +  CaH2O2=CaCl2  +  2 

The  gas  may  be  dried  by  being  passed  through  a  cylinder  con- 
taining lumps  of  quicklime,t  and  may  then  be  collected  either  by 
upward  displacement  or  in  the  mercurial  trough.  On  account  of 
its  extreme  solubility  it  cannot  be  collected  over  water. 

(3.)  Ammonia  is  formed  by  the  action  of  nascent  hydrogen  upon 
salts  of  nitrous  and  nitric  acid,  thus  — 


This  method  is  often  made  use  of  in  the  quantitative  estimation 
of  nitrates  in  drinking  water. 

(4.)  When  nitrogenous  organic  matter  is  subjected  to  destruc- 
tive distillation,  that  is,  strongly  heated  out  of  contact  with  air, 
ammonia  is  formed  ;  hence  when  coal,  which  usually  contains  about 
2  per  cent,  of  nitrogen,  is  distilled  in  the  process  of  the  manu- 
facture of  ordinary  illuminating  gas,  one  of  the  products  of  the 
decomposition  is  ammonia.  The  "ammoniacal  liquor"  of  the 
gas  works  is  the  source  of  all  ammonia  salts  at  the  present  day. 
The  liquor  is  boiled  with  milk  of  lime,  and  the  ammonia  thus 
expelled  is  absorbed  by  sulphuric  acid.  The  ammonium  sulphate 
so  obtained  is  purified  by  recrystallisation. 

Properties.  —  Ammonia  is  a  colourless  gas,  having  a  powerfully 
pungent  smell,  and  a  strong  caustic  taste.  It  is  lighter  than  air, 
its  density  being  o.  589  (air=  i).  Ammonia  possesses  the  property 
of  alkalinity  in  a  very  high  degree  ;  it  turns  red  litmus  blue,  and 
yellow  turmeric  brown.  The  gas  is  unable  to  support  combustion, 

*  A  solution  of  mercuric  iodide  in  potassium  iodide,  rendered  alkaline  with 
potassium  hydroxide. 

f  The  usual  desiccating  agents,  namely,  sulphuric  acid,  or  phosphorus 
pentoxide,  are  inadmissible  in  the  case  of  ammonia,  as  this  gas  at  once  unites 
with  such  compounds. 

S 


2/4 


Inorganic  Chemistry 


and  is  irrespirable.  Under  ordinary  conditions  ammonia  is  not 
combustible,  but  if  the  air  be  heated  or  if  the  amount  of  oxygen  be 
increased  the  gas  will  then  burn  with  a  flame  of  a  characteristic 
yellow-ochre  colour.  This  behaviour  of  ammonia  as  regards  com- 
bustibility is  most  conveniently 
illustrated  by  means  of  the  ap- 
paratus shown  in  Fig.  56.  A 
stream  of  the  gas  obtained  by 
gently  heating  a  quantity  of  the 
strong  aqueous  solution  in  a 
small  flask  is  delivered  through 
a  tube  which  is  surrounded  by 
a  wider  glass  tube.  Through 
the  cork  which  carries  this  tube 
a  second  tube  passes,  through 
which  a  supply  of  oxygen  can 
be  passed.  On  applying  a 
lighted  taper  to  the  jet  of  am- 
monia as  it  issues  from  the  tube 
it  will  be  noticed  that  the  gas 
burns  in  the  heated  air  round 
the  flame  of  the  taper,  but  is 
unable  to  continue  burning  when 
the  taper  is  withdrawn.  If  now 
a  gentle  stream  of  oxygen  be 
admitted  into  the  annular  space 
between  the  two  tubes  the  ammonia  readily  ignites,  and  continues 
to  burn  with  its  characteristic  flame.  On  cutting  off  the  supply 
of  oxygen  the  flame  of  the  burning  ammonia  languishes  and 
dies  out. 

Ammonia  is  extremely  soluble  in  water  ;  i  c.c.  of  water  at  o°  C., 
and  at  the  standard  pressure,  dissolves  1148  c.c.  of  ammonia, 
measured  at  o°  C.  and  760  mm.  The  solubility  rapidly  decreases 
as  the  temperature  rises,  as  will  be  seen  by  the  following  table  : — 


FIG.  56. 


i  c.c.  of  Water  at 
760  mm.  Dissolves 

At     0°   . 


i6c 


30 
50° 


Grammes,  NHiV 

.  0.875  • 

.  0.713  . 

.  0.582  . 

.  0.403  . 

.  0.229  . 


c.c.  at  o°  C.  and 
760  mm. 

.  1148 
•  923 
.  764 
.  529 
106 


Ammonia 


275 


When  a  solution  of  ammonia  is  heated  the  gas  is  rapidly  evolved, 
and  at  the  boiling  temperature  the  whole  of  it  is  given  up. 

The  great  solubility  of  this  gas  in  water  may  be  shown  by  filling 
a  large  bolt-head  flask  with  ammonia  by  displacement,  the  flask 
being  closed  by  means  of  a  cork  through  which  a  long  tube  passes, 
as  shown  in  Fig.  57.  On  removing  the  cork  from  the  end  of  the 
tube  water  slowly  rises  until  it  reaches  the  top,  and  as  soon  as  the 
first  drops  enter  the  globe  the  absorption 
proceeds  with  great  rapidity,  the  water  being 
forced  up  the  tube  in  the  form  of  a  fountain, 
which  continues  until  the  flask  is  filled. 

Commercial  liqtior  armnonicB  is  prepared 
by  passing  ammonia  gas  into  water  ;  the 
strongest  solution  has  a  specific  gravity  of 
0.882  at  15°,  and  contains  35  per  cent,  of 
ammonia.  During  the  process  of  solution 
heat  is  liberated,  and  when  the  gas  is  again 
expelled  the  same  amount  of  heat  is  reab- 
sorbed.  If  a  rapid  stream  of  air  be  driven 
through  a  quantity  of  strong  ammonia  solu- 
tion, contained  in  a  glass  flask,  the  ammonia 
gas  is  quickly  expelled  ;  and  if  the  flask 
be  placed  upon  a  wooden  block,  as  seen  in  ' 
Fig.  58,  upon  which  a  few  drops  of  water 
have  been  poured,  it  will  be  found  that  after 
a  few  moments  the  flask  will  have  become 

firmly  frozen  to  the  block.     By  the  rapid  evaporation  of  ammonia 
in  this  way  it  is  possible  to  lower  the  temperature  to  —40°  C. 

Ammonia  is  an  easily  liquefiable  gas  ;  thus  at  15.5°  it  requires  a 
pressure  of  6.9  atmospheres,  and  at  o°  only  4.2  atmospheres,  in  order 
to  liquefy  it.  The  gas  was  first  liquefied  by  Faraday  (1823)  by  heat- 
ing in  one  limb  of  a  closed  and  bent  glass  tube  (see  Fig.  2)  a  quantity 
of  a  compound  of  ammonia  with  silver  chloride,  the  other  limb  of 
the  tube  being  immersed  in  a  freezing-mixture.  The  experiment 
may  be  made  in  a  tube  constructed  as  seen  in  Fig.  59.  The  wide 
limb  is  nearly  filled  with  dry  precipitated  silver  chloride  which  has 
been  saturated  with  ammonia  gas.  This  compound  melts  at 
about  38°,  and  at  a  somewhat  higher  temperature  it  gives  up 
its  ammonia.  If  the  narrow  limb  of  the  tube  be  immersed  in  a 
freezing-mixture  while  the  compound  is  being  heated,  the  com- 
bined influence  of  the  cold  and  the  pressure  exerted  by  the 


FIG.  57. 


276  Inorganic  Chemistry 

evolved  ammonia  will  cause  the  gas  to  liquefy  and  collect  in  the 
cold  portion  of  the  tube.  On  removing  the  tube  from  the  freezing- 
mixture  and  allowing  the  other  end  to  cool,  the  liquid  ammonia 
will  boil  off  and  be  reabsorbed  by  the  silver  chloride,  reforming 
the  original  compound. 

Liquid  ammonia  is  easily  obtained  in  larger  quantity  by  passing 
the  gas  through  a  glass  tube  immersed  in  a  bath  of  solid  carbonic 
acid  and  ether.  Liquid  ammonia  is  a  colourless,  mobile,  and 
highly  refracting  liquid,  boiling  at  —33.7°,  and  having  a  specific 
gravity  at  o°  of  0.6234.  When  cooled  below  —75°  it  solidifies  to  a 
mass  of  white  crystals. 

Liquid  ammonia  dissolves  the  metals  sodium  and  potassium,  the 
solution  in  each  case  being  of  an  intense  blue  colour.  On  the 
evaporation  of  the  liquid  the  metal  is  deposited  unchanged. 


FIG.  58.  FIG.  59. 

During  the  evaporation  of  liquid  ammonia,  boiling  as  it  does  al- 
so low  a  temperature  as  —  33.7°,  a  rapid  absorption  of  heat  takes 
place,  and  as  this  substance  is  so  easily  obtained  it  was  one  of  the 
earliest  liquids  employed  for  the  artificial  production  of  ice.  Various 
ice-making  machines  have  been  invented  by  M.  Carrd,  in  which  the 
reduction  of  temperature  required  is  obtained  by  the  evaporation 
of  liquid  ammonia. 

Ammonia  is  decomposed  into  its  elements  at  a  temperature 
below  a  red  heat.  In  this  decomposition  two  volumes  of  ammonia 
give  one  volume  of  nitrogen  and  three  volumes  of  hydrogen.  The 
gaseous  products,  therefore,  obtained  by  passing  ammonia  through 
a  red-hot  tube  are  inflammable.  In  the  same  way,  when  electric 
sparks  are  passed  through  ammonia,  the  gas  is  resolved  into  its 
constituents.  By  performing  this  experiment  upon  a  measured 
volume  of  ammonia  confined  in  a  eudiometer  over  mercury,  it  will 
be  found  that  after  the  passage  of  the  sparks  for  a  short  time  and 


Ammonia 


277 


the  readjustment  of  the  levels  of  mercury,  the  original  volume  of 
the  gas  has  been  doubled. 

The  fact  that  the  hydrogen  and  nitrogen  are  present  in  ammonia 
in  the  proportion  of  three  volumes  of  hydrogen  to  one  of  nitrogen 
can  be  shown  by  taking  advantage  of  the  fact  that  ammonia  is 
decomposed  by  chlorine,  the  latter  combining  with  the  hydrogen 
to  form  hydrochloric  acid  and  the  nitrogen  being  set  free.  This  is 
effected  by  means  of  the  apparatus  shown  in  Fig.  60.  The  long 
glass  tube,  divided  into  three  equal  divi- 
sions, is  filled  with  chlorine  and  closed  by 
a  cork  carrying  a  small  dropping  funnel.  A 
few  cubic  centimetres  of  strong  aqueous 
ammonia  are  poured  into  the  funnel  and 
allowed  to  enter  the  tube  drop  by  drop. 
As  the  first  two  or  three  drops  fall  into 
the  chlorine  it  will  be  seen  that  the  com- 
bination is  attended  with  a  feeble  flash  of 
light,  and  fumes  of  ammonium  chloride  are 
formed.  When  the  reaction  is  complete 
the  whole  of  the  chlorine  will  have  com- 
bined with  hydrogen  derived  from  the 
ammonia  to  form  hydrochloric  acid,  and 
this  in  its  turn  will  combine  with  the  excess 
of  ammonia  added,  forming  ammonium 
chloride.  This  substance  dissolves  in  the 
water.  A  small  quantity  of  dilute  sulphuric 
acid  is  next  introduced  by  means  of  the 
dropping  funnel  in  order  to  absorb  the 
remaining  excess  of  ammonia.  The  at- 
mospheric pressure  is  then  once  more  re- 
stored by  attaching  to  the  funnel  a  bent 
tube,  dipping  into  a  beaker  of  water,  as 

shown  in  the  figure,  and  when  the  water  is  allowed  to  enter 
it  will  be  found  to  flow  into  the  tube  until  it  reaches  the  second 
graduation.  The  gas  which  is  left  and  which  occupies  one  of  the 
divisions  of  the  tube  is  found  on  examination  to  be  nitrogen. 
This  one  measure  of  nitrogen,  therefore,  has  been  eliminated  from 
that  amount  of  ammonia  which  has  been  decomposed  by  the 
chlorine  with  which  the  tube  was  originally  filled.  Now  chlorine 
combines  with  its  own  volume  of  hydrogen,  therefore  the  volume 
of  hydrogen  which  was  in  combination  with  the  one  measure  of 


FIG.  60. 


278  Inorganic  Chemistry 

nitrogen  is  equal  to  the  volume  of  chlorine  contained  in  the  tube, 
that  is  to  say,  it  was  three  measures.  We  have,  therefore,  one 
measure  of  nitrogen  and  three  measures  of  hydrogen,  or,  in  other 
words,  ammonia  is  a  combination  of  nitrogen  and  hydrogen  in  the 
proportion  of  one  volume  of  nitrogen  to  three  volumes  of  hydrogen. 

In  contact  with  many  metals  at  a  moderately  high  temperature 
ammonia  is  decomposed  into  its  elements,  and  a  compound  of 
the  metal  with  nitrogen  is  formed.  In  this  way,  at  temperatures 
ranging  between  about  400°  and  800°  a  number  of  metallic  nitrides 
have  been  obtained.*  These  compounds  are  produced  by  passing 
a  rapid  stream  of  ammonia  gas  through  heated  porcelain  tubes 
containing  the  metal  in  the  form  of  either  wire,  foil,  or  fine  powder. 
When  heated  in  an  atmosphere  of  hydrogen,  these  nitrides  are 
decomposed  into  nitrogen  and  the  respective  metal,  hence  they 
can  only  be  produced  in  the  presence  of  a  large  excess  of  am- 
monia gas. 

Ammonia  combines  directly  with  acids  forming  salts,  known 
as  ammonium  salts,  in  which  the  nitrogen  functions  as  a  pentad 
element ;  thus  with  hydrochloric  and  sulphuric  acids  it  forms  respec- 
tively ammonium  chloride  and  ammonium  sulphate — 

NH3  +  HC1  =  (NH4)C1. 
2NH3  +  H2S04=(NH4)2S04. 

(The  ammonium  salts  will  be  described  with  the  compounds  of 
the  alkali  metals.) 

Hydrazine  (diamidogen],  NH2' NH2  or  N2H4.— This  compound  was  first 
prepared  by  Curtius  (1887).  It  is  obtained  from  a  salt  of  an  organic  acid 

N\ 
known  as  diazo-acetic  acid,   ||-      CH'COOH.     When  the  ethereal  salt  of  this 

N/ 

acid  is  acted  upon  by  potassium  hydroxide,  the  potassium  salt  of  another 
acid  is  formed,  namely  triazo-acetic  acid.  This  we  may  regard  as  merely  a 
polymer  of  the  first  acid,  and  represent  its  formula  (N2 :  CH'COOH)3. 
When  this  compound  is  digested  with  dilute  sulphuric  acid  it  is  converted 
into  hydrazine  sulphate  and  oxalic  acid.  Thus,  employing  the  simple  formula 
for  the  acid — 

N2 :  CH  -COOH  +  H2SO4 + 2H2O = N2H4H2SO4  +  H2C2O4. 

Hydrazine  may  also  be  prepared  from  purely  inorganic  sources.  When 
hydrogen  potassium  sulphite  is  acted  upon  by  potassium  nitrite,  a  compound 
known  as  potassium  dinitroso-sulphonate  is  produced,  O :  N'N-OK'KSO3. 

*  Beilby  and  Henderson,  Jour.  Chem.  Soc. ,  November  1901. 


Hydrazine  279 

The  mechanism  of  the  reaction  will  be  made  clearer  if  the  formula  for  the 
nitrite  be  written  O  :  N'OK.  Thus — 

20 :  N'OK  +  2HKSO3=O :  N-N'OK-KSO3+K2SO4. 

By  the  action  of  nascent  hydrogen  (from  sodium  amalgam)  at  the  temperature 
of  ice,  this  compound  is  converted  into  the  potassium  salt  of  hydrazine 
sulphonate — 

O:N-N-OK-KSO3+6H  =  H2N-NH-KSO3+KHO  +  H2O. 

And  this  compound  on  distillation  with  potassium  hydroxide  yields  hydrazine — 
H2N-NH-KSO3+HKO=H2N-NH2+K2SO4. 

The  base  itself  may  also  be  obtained  by  heating  together  in  a  sealed  tube,  to 
a  temperature  of  170°,  hydrazine  hydrate,  N2H4,H2O,  and  barium  monoxide. 
Under  these  circumstances  the  barium  oxide  takes  up  the  water  from  the 
hydrazine  hydrate,  according  to  the  equation — 

BaO  +  N2H4H20 = Ba(  HO)2 + N2H4. 

When  the  tube  is  opened,  the  gaseous  hydrazine,  which  is  under  considerable 
pressure,  rushes  out  of  the  tube,  forming  dense  fumes  in  contact  with  the 
atmospheric  moisture,  with  which  it  combines  with  great  readiness. 

Hydrazine  Hydrate,  N2H4H2O. — The  compound  formed  by  the  combina- 
tion of  hydrazine  with  water  is  obtained  by  distilling  hydrazine  sulphate, 
N2H4H2SO4,  with  an  aqueous  solution  of  potassium  hydroxide  (caustic  potash) 
in  a  vessel  of  silver.  It  is  a  colourless,  fuming,  powerfully  corrosive  liquid, 
which  boils  at  118.5°.  ^  attacks  glass,  cork,  and  indiarubber,  and  can  only 
be  prepared  in  vessels  of  silver  or  platinum  which  are  screwed  together  at  their 
junctions.  With  the  halogen  acids  it  forms  two  series  of  salts,  in  which  either 
one  or  two  molecules  of  the  halogen  acid  are  present :  thus  with  hydrochloric 
acid  we  have — 

Hydrazine  monohydrochloride      .        .         .     N2H4HC1. 
Hydrazine  dihydrochloride  ....     N2H42HC1. 

Hydrazine  and  its  salts  act  as  powerful  reducing  agents,  and  give  the  charac- 
teristic red  precipitate  of  cuprous  oxide  when  added  to  Fe'hling's  solution. 
This  reaction  serves  to  immediately  distinguish  these  compounds  from 
ammonium  salts. 

/N 
Hydrazoic  Acid  or  Azoimide,  HN3  or  HN      ||  .—Discovered  by  Curtius 

(1890).  The  sodium  salt  is  prepared  by  boiling  benzoylazo-imide  with  sodium 
hydroxide,  when  sodium  benzoate  and  sodium  hydrazoate  are  formed,  thus — 

/N  /N 

C6H5CO-N      ||  +  2NaHO=C6H5CO-ONa+Na-N      ||  +HaO. 


280  Inorganic  Chemistry 

It  is  also  produced  when   sodamide  (obtained  by  heating  sodium  in  dry 
ammonia  gas)  is  heated  to  200°  in  a  stream  of  nitrous  oxide  *  — 

2NH2Na+N20=NaN3+NaHO  +  NH3. 

The  sodium  hydrazoate  so  obtained  is  then  gently  warmed  with  dilute  sulphuric 
acid,  when  sodium  sulphate  and  hydrazoic  acid  are  formed,  thus  — 


Properties.  —  This  compound  is  a  colourless  volatile  liquid,  boiling  at  37°. 
The  vapour  possesses  a  most  unpleasant  and  powerfully  penetrating  odour. 
If  inhaled,  even  when  largely  diluted  with  air,  it  exerts  an  irritating  action 
upon  the  mucous  membrane.  As  its  name  denotes,  it  is  an  acid  substance, 
and  in  many  of  its  properties  it  strongly  resembles  the  halogen  acids.  The 
compound  is  extremely  soluble  in  water,  and  forms  a  strongly  acid  liquid 
which  smells  of  the  vapour.  This  solution  when  boiled,  finally  assumes  a 
definite  strength,  and  yields  on  distillation  an  aqueous  acid  of  constant  com- 
position, in  this  respect  resembling  aqueous  hydrochloric  acid,  q.v. 

In  its  constitution  this  acid  may  be  compared  with  hydrocyanic  acid,  and 
with  the  halogen  acids  — 

H(N3);  H(CN);  H(C1)  ;  H(Br), 

in  which  the  radical  cyanogen  (CN),  or  the  halogen  elements,  Cl  and  Br,  are 
replaced  by  the  group  consisting  of  three  nitrogen  atoms. 

When  a  solution  of  hydrazoic  acid  is  added  to  a  solution  of  silver  nitrate,  a 
white  precipitate  of  silver  hydrazoate  is  formed,  strongly  resembling  silver 
cyanide  and  silver  chloride  in  appearance.  This  silver  salt,  however,  is  not 
acted  upon  by  light  in  the  way  the  chloride  is,  and  it  differs  also  in  being 
extremely  explosive.  A  minute  quantity  of  the  compound,  when  touched  with 
a  hot  wire,  detonates  violently. 

This  instability  and  tendency  to  explode  is  characteristic  v  of  the  acid  and 
most  of  its  salts.  The  sodium  salt,  however,  may  be  heated  to  about  250° 
before  it  decomposes. 

When  gaseous  hydrazoic  acid  is  mixed  with  gaseous  ammonia,  dense  white 
fumes  are  formed,  consisting  of  ammonium  hydrazoate.  These  two  hydrides 
of  nitrogen,  apparently  so  similar,  but  in  reality  so  widely  different,  unite  to 
form  the  ammonium  salt,  just  as  gaseous  hydrochloric  acid  and  ammonia 
combine  to  form  ammonium  chloride,  thus  — 

NH3+H(N3)  =  NH4(N3). 
NH3  +  HC1=NH4C1. 

The  alkaline  hydride  of  nitrogen,  ammonia,  combines  with  the  acid  hydride  of 
nitrogen,  hydrazoic  acid,  and  forms  the  salt  ammonium  hydrazoate  NH4N3 
or  N4H4.  The  salts  of  this  acid  are  sometimes  called  nitrides,  thus  sodium 
nitride,  NaNs. 

*  See  Experiment  298,  "Chemical  Lecture  Experiments,"  new  ed. 


Hydroxylamine  281 

HYDROXYLAMINE. 

Formula,  NH2(OH). 
Discovered  by  Lessen  in  1865. 

Modes  of  Formation.—  (  i.)  By  the  action  of  nascent  hydrogen 
upon  nitric  oxide,  nitric  acid,  or  certain  nitrates  — 

2NO  +  3H9  =  2NH2(OH). 
HNO3  +  3H2  =  2H20  +  NH2(OH). 

The  nascent  hydrogen  is  evolved  from  tin  and  hydrochloric  acid, 
and  a  stream  of  nitric  oxide  passed  through  the  mixture.  The 
hydrochloride  of  hydroxylamine  is  thus  obtained.  The  dissolved 
tin  is  then  precipitated  as  sulphide  by  means  of  hydrogen  sulphide, 
the  filtered  liquid  evaporated  to  dryness,  and  the  hydroxylamine 
hydrochloride  dissolved  out  of  the  residue  with  absolute  alcohol. 
On  evaporating  this  solution  the  salt  is  obtained  in  the  form  of 
white  crystals. 

(2.)  By  the  interaction  of  alkali  nitrites  and  metasulphites,  and 
the  subsequent  prolonged  boiling  of  the  hydroxylamine  disulphonate 
so  obtained  with  water  —  - 

2KN  02  +  3K2S03,SO2  +  H2O  =  2K,SO3  +  2N(OH)(SO3K)2. 


The  potassium  sulphate  is  separated  from  the  hydroxylamine 
sulphate  by  crystallisation. 

Hydroxylamine  itself,  in  aqueous  solution,  may  be  obtained  from 
the  sulphate  by  the  addition  of  baryta  water.  In  the  anhydrous 
state  it  is  produced  by  the  distillation  oT  hydroxylamine  phosphate 
under  reduced  pressure,  the  distillate  being  solidified  by  immersion 
in  ice.  If  this  is  dissolved  in  absolute  alcohol  and  the  solution 
cooled  to  about  —  18°  the  pure  compound  separates  out  in  white 
scales  or  needles. 

Properties.  —  Hydroxylamine  melts  at  33°,  and  under  reduced 
pressure  may  be  boiled  and  distilled  ;  but  although  tolerably  stable 
at  the  ordinary  temperature  it  decomposes  with  explosion  when 
heated  above  about  100°  at  atmospheric  pressure.  It  is  readily 
soluble  in  water  yielding  a  strongly  alkaline  solution,  which  pre- 
cipitates silver  and  mercury  from  solutions  of  their  salts,  and 
which  reduces  cupric  salts  with  precipitation  of  red  cuprous  oxide. 

Hydroxylamine  is  a  base,  and  may  be  regarded  as  ammonia,  in 
which  one  of  the  hydrogen  atoms  has  been  replaced  by  the  monad 
group  hydroxyl  (OH).  Like  ammonia  it  unites  with  acids  forming 
salts,  without  the  elimination  of  water. 

NH,OH  +  HCl  =  NH3OHCl(orNH2OH,HCl). 
2NH2OH  +  H2S04  =  (NH3OH)2S04  (or  2NH2OH,H2SO4). 

The  salts  of  hydroxylamine  all  decompose  on  the  application  of 


282  Inorganic  Chemistry 

heat,  with  a  more  or  less  sudden  and  violent  evolution  of  gas  ;  thus 
the  nitrate  breaks  up  with  almost  explosive  violence  into  nitric 
oxide  and  water — 

NH2OH'HNO3  =  2NO  +  2H2O. 

AMMON-SULPHONATES. 

These  compounds  may  be  regarded  as  derived  from  ammonia,  by  the 
gradual  replacement  of  the  hydrogen  by  the  group  SO3H  or  SO2OH. 

Ammon-sulphonic  acid      .         .         .     NH2(SO3H). 
Ammon-disulphonic  acid  .         .     NH(SO3H)2. 

Ammon-trisulphonic  acid          .         .     N(SO3H)3. 

Potassium  ammon-trisulphonate  is  precipitated  as  a  crystalline  salt  when 
excess  of  a  solution  of  potassium  sulphite  is  added  to  a  solution  of  potassium 
nitrite — 

3K2SO3  +  KNO2+2H2O==4KHO  +  N(SO3K)3. 

Prolonged  boiling  with  water  converts  it  first  into  the  ammon-disulphonate — 

N(SO3K)3+  H2O=NH(SO3K)2  +  H  KSO4, 
and  finally  into  ammon-sulphonate— 

NH(S03K)2+  H20=NHo(S03K)  +  HKSO4. 

Ammon-sulphonic  acid  is  a  stable  crystalline  body  ;  the  other  two  acids  are 
only  known  in  their  salts. 

When  an  ice-cold  solution  of  sodium  nitrite  is  added  to  hydrogen  sodium 
sulphite,  a  compound  is  obtained  which  may  be  regarded  as  derived  from 
ammon-trisulphonic  acid  by  the  replacement  of  one  of  the  groups,  SO3H,  by 
hydroxyl,  OH— 

NaN02+2NaHS03=N(OH)(S03Na)2  +  NaHO. 

On  the  addition  of  a  saturated  solution  of  potassium  chloride  in  the  cold,  the 
sodium  salt  is  converted  into  the  potassium  salt,  which  slowly  crystallises  from 
the  solution,  with  two  molecules  of  water,  N(OH)(SO3K)2,2H2O. 

This  potassium  hydroxylamine  disulphonate  is  an  unstable  compound,  and 
on  boiling  with  water  the  two  SO3K  groups  are  replaced  by  hydrogen,  forming 
first  potassium  hydroxylamine  monosulphonate,  NH(OH)SO3K,  and  finally 
hydroxylamine,  NH2OH. 

COMPOUNDS  OF  NITROGEN  WITH  THE  HALOGEN  ELEMENTS. 

Nitrogen  Trichloride,  NC13. — This  compound  was  discovered  by  Dulong 
(1811).  Its  true  composition  was  proved  by  Gattermann  (1888). 

Mode  of  Formation.— Nitrogen  trichloride  is  obtained  by  the  action  of 
chlorine  upon  ammonium  chloride— 

NH4Cl+3Cla=NCl8+4HCl. 


Nitrogen  Iodide  283 

When  a  solution  of  ammonium  chloride  is  electrolysed,  the  chlorine,  which  is 
evolved  at  the  positive  electrode,  acts  upon  the  ammonium  chloride,  forming  the 
trichloride  of  nitrogen.* 

Properties.  —  Nitrogen  trichloride  is  a  thin  oily  liquid,  of  a  pale-yellow 
colour,  and  having  a  specific  gravity  of  1.65.  It  is  very  volatile,  and  has  an 
unpleasant  pungent  smell,  the  vapour  being  extremely  irritating  to  the  eyes.  It 
is  the  most  dangerously  explosive  compound  known,  and  when  suddenly 
heated,  or  brought  into  contact  with  grease,  turpentine,  or  phosphorus,  it  at  once 
explodes.  It  also  explodes  on  exposure  to  sunlight.  At  a  temperature  of  71° 
it  may  be  distilled,  but  the  operation  is  one  of  the  utmost  danger.  Nitrogen 
trichloride  is  decomposed  by  ammonia,  forming  ammonium  chloride  and  free 
nitrogen  ;  hence  in  the  preparation  of  nitrogen  by  the  action  of  chlorine  upon 
ammonia,  the  presence  of  an  excess  of  ammonia  prevents  the  formation  of  this 
dangerous  compound. 

Nitrogen  Tribr  oroide,  NBr3.  —  When  potassium  bromide  is  added  to  nitro- 
gen trichloride  beneath  water,  a  red,  oily,  hignly  explosive  substance  is 
obtained,  believed  to  be  the  tribromide  of  nitrogen. 

Nitrogen  Iodide,  N2H3I3.  —  When  strong  aqueous  ammonia  is  added  to 
powdered  iodine,  a  brown-coloured  powder  is  formed  which  has  violently 
explosive  properties.  Also  when  alcoholic  solutions  of  iodine  and  of  ammonia 
are  mixed,  a  brown  and  highly  explosive  compound  is  produced. 

Curtois,  who  first  prepared  the  substance,  believed  it  to  have  the  composi- 
tion NI3,  and  this  view  was  held  by  Gay-Lussac  and  others.  Gladstone  and 
others  considered  that  the  substance  contained  one  atom  of  hydrogen,  and  that 
the  formula  NHI2  expressed  the  composition.  The  investigations  of  Szuhay 
(1893)  a^so  ted  him  to  believe  that  the  compound  obtained  by  the  addition 
of  an  excess  of  aqueous  ammonia  to  a  concentrated  solution  of  iodine  in 
potassium  iodide  has  the  composition  NHI2. 

The  subject  has  recently  been  reinvestigated  by  Chattaway  (Proc.  Chem.  Soc., 
1899),  who  for  the  first  time  appears  to  have  obtained  the  compound  in  a  state 
of  purity  by  the  addition  of  ammonia  to  a  solution  of  potassium  hypoiodite. 
Under  these  circumstances  the  substance  separates  out  in  the  form  of  well- 
defined  crystals  having  a  composition  expressed  by  the  formula  N2H3I3,  which 
may  be  regarded  either  as  NI3,NH3  or  NHI2,NH2I.  The  equations  represent- 
ing the  formation  of  the  compound  may  be  thus  expressed  —  • 

KIO  +  NH4HO=NH4IO+  KHO, 

3NH4IO=N2H3I3+NH3+3H2O. 

The  reaction  which  takes  place  when  the  compound  is  obtained  by  the 
action  of  iodine  upon  strong  ammonia  appears  also  to  involve  the  first  forma- 
tion of  the  unstable  ammonium  hypoiodite,  thus  — 


I2  +  2NH4HO=NH4 

which  then  breaks  up  as  shown  above. 

Properties.  —  Nitrogen  iodide  is  a  copper-coloured  glittering  crystalline  com- 
pound, appearing  red  by  transmitted  light.  In  the  amorphous  state,  as  obtained 
by  the  action  of  iodine  upon  strong  ammonia,  it  presents  the  appearance  of 

*  See  "Chemical  Lecture  Experiments,"  new  ed.,  No.  301. 


284  Inorganic  Chemistry 

a  dark  chocolate-brown  powder.  When  moist  it  may  be  handled  without  much 
risk  of  explosion,  although  it  has  been  known  to  explode  even  under  water. 
When  dry  the  substance  is  extremely  explosive  ;  the  shock  caused  by  the  tread 
of  a  fly  upon  it  is  more  than  sufficient  to  explode  it ;  even  falling  dust  particles 
will  sometimes  cause  it  to  explode. 

When  nitrogen  iodide  is  placed  in  dilute  aqueous  ammonia,  and  exposed  to 
bright  light,  it  is  decomposed,  and  bubbles  of  nitrogen  are  seen  escaping  from 
the  compound — 

N2H3I3=N2+3HI, 

the  hydriodic  acid  being  neutralised  by  the  ammonia  present.  At  the  same 
time  a  small  quantity  of  the  compound  is  converted  into  ammonium  hypoiodite, 
which  being  unstable  slowly  passes  into  the  iodate,  thus — 

N2H3I3+3H2O  +  NH3=3NH4IO, 
3NHJO=NH4I03  +  2NH4I. 


CHAPTER   VIII 
CARBON 

Symbol,  C.     Atomic  weight  =  12.00. 

Occurrence. — This  element  is  capable  of  assuming  three  allo 
tropic  forms,  and  it  occurs  free  in  nature  in  each  of  these  modifica- 
tions, viz.,  diamond,  graphite,  and  charcoal. 

In  combination  with  oxygen,  carbon  occurs  in  carbon  dioxide,  a 
gas  which  is  present  in  the  air,  being  a  constant  product  of  com- 
bustion and  respiration.  In  combination  with  hydrogen  it  occurs 
as  marsh  gas.  Carbon  is  a  constituent  of  all  the  natural  car- 
bonates, such  as  limestone,  dolomite,  &c.,  which  form  an  important 
fraction  of  the  earth's  crust,  and  it  is  also  an  essential  constituent 
of  all  organic  substances. 

DIAMOND. 

Occurrence. — This  substance  has  been  known  and  prized  frorr. 
the  remotest  antiquity.  It  is  found  in  various  parts  of  India, 
mostly  in  river  gravels  and  superficial  deposits,  in  Brazil,  South 
Africa,  Australia,  and  various  parts  of  the  United  States.  The 
diamond  has  also  recently  been  obtained  from  extra-terrestrial 
sources.  In  a  meteorite  which  fell  in  Russia  on  September  22, 
1886,  carbon  was  found,  partly  as  amorphous  and  partly  as  ada- 
mantine carbon. 

The  diamond  form  of  carbon  is  found  of  various  colours  ;  some- 
times it  is  dark  grey,  or  even  black,  stones  of  these  colours  being 
known  as  carbonado  and  bort.  The  former  of  these  is  extremely 
hard,  and  is  of  great  value  for  use  in  rock-boring  and  drilling 
instruments.  Bort  is  used  in  the  crushed  condition  by  lapidaries 
for  grinding  and  polishing. 

Occasionally  the  diamond  is  found  coloured  blue,  or  red,  or 
green  by  traces  of  foreign  materials.  Some  of  these  coloured 


286  Inorganic  Chemistry 

stones  are  of  great  value  as  gems:  the  well-known  "Hope" 
diamond,  a  stone  weighing  44^  carats,  has  a  fine  sapphire 
colour. 

The  origin  of  the  diamond  is  unknown,  although  many  theories 
have  been  put  forward  to  explain  its  formation.  Newton's  famous 
suggestion,  that  diamond  was  "  an  unctuous  substance  coagulated," 
was  based  upon  its  remarkably  high  refractive  index.  The  cellular 
structure  which  is  sometimes  to  be  seen  in  the  ash  that  is  left  when 
the  diamond  is  burnt  seems  to  indicate  that  it  is  of  vegetable 
origin. 

Modes  of  Formation. — Innumerable  attempts  have  been  made 
to  effect  the  crystallisation  of  carbon  in  the  adamantine  form  ;  but 
while  it  is  readily  possible  to  convert  this  variety  of  carbon  into  its 
allotropes  graphite  and  charcoal,  the  transformation  of  these  back 
again  to  the  diamond  is  a  problem  that  is  beset  with  the  greatest 
difficulties.  Moissan  has  recently  shown  *  that  the  carbon,  which 
is  capable  of  being  dissolved  in  molten  iron,  and  which  is  usually 
deposited  in  the  graphitic  form  on  cooling,  can,  under  certain 
conditions,  be  caused  to  take  up  the  adamantine  form. 

By  raising  the  temperature  of  the  iron  to  about  3000°  by  means 
of  an  electric  furnace,  and  then  suddenly  cooling  the  molten  mass 
by  plunging  the  crucible  into  water  or  molten  lead,  until  the  cooled 
and  solidified  surface  is  at  a  dull  red  heat,  an  enormous  pressure  is 
brought  to  bear  upon  the  interior  and  still  liquid  portion.  Under 
these  circumstances,  a  part  of  the  carbon  which  is  deposited  by  the 
slowly  cooling  mass  was  found  by  Moissan  to  be  in  the  adamantine 
form.  On  dissolving  the  iron  in  hydrochloric  acid,  amongst  the 
carbonaceous  residue  were  found  fragments  having  a  specific 
gravity  between  3.0  and  3.5,  and  sufficiently  hard  to  scratch  ruby. 
Some  of  the  fragments  were  the  black  or  carbonado  variety,  while 
others  were  transparent.  On  combustion  in  oxygen,  Moissan 
proved  that  these  were  really  carbon  in  the  diamond  form. 

Properties. — The  diamond  in  its  purest  condition  is  a  colourless 
crystalline  substance.  Its  crystalline  forms  belong  to  the  cubic 
system,  and  appear  to  some  extent  to  be  characteristic  of  the 
locality  in  which  the  element  occurs.  It  is  extremely  hard  and 
moderately  brittle.  When  struck  with  a  hammer  the  diamond  not 
only  splits  along  its  cleavage-planes,  but  also  in  other  directions, 
with  a  conchoidal  fracture.  It  does  not  conduct  electricity.  The 

*  Comptes  Rendus  de  V  Academic  des  Sciences,  vol.  cxvi.  p.  218. 


Carbon 


287 


specific  gravity  of  diamond  varies  slightly  in  different  specimens, 
the  mean  being  about  3.5.  Its  refractive  index  is  higher  than  that 
of  any  other  substance,  and  it  is  this  property  which  gives  its 
peculiar  beauty  and  brilliancy  to  the  cut  stone. 

The  value  of  diamond  as  a  gem  depends  largely  upon  its  colour- 
lessness, except  in  the  case  of  those  rare  instances  where  the 
colour  is  quite  definite  and  also  pleasing,  such  as  distinct  red,  blue, 
or  green. 

When  diamond  is  strongly  heated  it  becomes  black,  and  in- 
creases in  bulk,  being  converted  into  a  substance  having  the 
properties  of  coke.  Lavoisier  (1772)  was  the  first  to  show  that 
the  diamond  was  a  combustible  body,  and  that  it  yielded  carbon 
dioxide.  Davy  (1814)  showed  that  carbon  dioxide  was  the  only 
product  of  its  combustion,  and  proved  that 
diamond  was  pure  carbon. 

The  combustion  of  diamond  in  oxygen  may 
readily  be  accomplished  by  means  of  the  ap- 
paratus shown  in  Fig.  61.  A  fragment  of 
diamond  is  supported  upon  a  small  gutter  of 
platinum  foil,  which  bridges  across  two  stout 
copper  wires,  A.  These  wires  pass  through  a 
cork  in  a  perforated  glass  plate,  and  are 
lowered  into  a  cylinder  of  oxygen.  By  the 
passage  of  an  electric  current  the  little  plati- 
num boat  can  be  strongly  heated,  when  the 
diamond  will  become  ignited,  and  continue  to 
burn  brilliantly  in  the  oxygen,  with  the  formation  of  carbon  dioxide. 
The  ash,  which  is  always  left  after  a  diamond  has  been  burnt, 
varies  from  0.2  to  0.05  per  cent,  of  the  stone.  It  is  found  usually 
to  contain  ferric  oxide  and  silica. 


FIG.  61. 


GRAPHITE. 

Occurrence. — This  second  allotrope  of  carbon  is  much  more 
plentiful  in  nature  than  the  first.  It  is  found  in  large  quantities 
in  Siberia,  Ceylon,  and  various  parts  of  India.  In  England  the 
chief  source  of  graphite  has  been  the  mines  at  Borrowdale  in 
Cumberland  ;  this  supply  is  now  practically  exhausted.  Enor- 
mous quantities  of  very  pure  graphite  are  now  obtained  from  the 
(Eureka  Black-Lead  Mines  in  California.  Graphite  also  occurs 
in  many  specimens  of  meteoric  iron. 


288  Inorganic  Chemistry 

Mode  of  Formation. — Molten  iron,  especially  when  it  contains 
silicon,  is  capable  of  dissolving  a  considerable  amount  of  carbon, 
which,  on  cooling,  is  deposited  in  the  form  of  black  shining  crystals 
of  graphite.  Occasionally  considerable  quantities  of  graphite  are 
found  deposited  in  this  way  in  iron-smelting  furnaces,  to  which  the 
name  "  kish  K  has  been  applied. 

Graphite  is  now  manufactured  by  heating  a  mixture  of  97  parts 
of  amorphous  carbon  (charcoal  or  coke)  and  3  parts  of  iron  in  an 
electric  furnace.  It  was  formerly  believed  that  at  the  high  tem- 
perature of  the  electric  arc  amorphous  carbon  was  converted 
directly  into  the  graphitic  modification  ;  but  it  has  recently  been 
shown  (Acheson)  that  pure  charcoal  does  not  by  itself  undergo 
this  transformation  ;  that  the  change,  in  reality,  takes  place 
through  the  intermediate  formation  of  a  metallic  carbide.  The 
product  obtained  is  practically  free  from  iron,  as  the  metal  is 
volatilised  at  the  high  temperature. 

Properties. — Graphite  is  a  soft,  shiny,  greyish-black  substance, 
which  is  smooth  and  soapy  to  the  touch.  It  is  usually  found  in 
compact  laminated  masses,  but  sometimes  crystallised  in  six-sided 
plates.  Its  specific  gravity  varies  in  different  specimens,  aver- 
aging about  2.5.  Graphite  is  a  good  conductor  of  both  heat  and 
electricity. 

When  strongly  heated  in  oxygen,  graphite  takes  fire  and  burns, 
forming  carbon  dioxide,  and  leaving  an  ash  consisting  of  silica, 
alumina,  and  oxide  of  iron.  Graphite  has  been  found  by  Regnault 
to  contain,  usually,  traces  of  hydrogen.  Graphite  is  employed  for 
the  manufacture  of  ordinary  lead  pencils  ;  for,  on  account  of  its 
softness,  it  leaves  a  black  mark  upon  paper  when  drawn  across  it. 
For  the  purposes  of  the  pencil  manufacture  the  natural  graphite  is 
ground  to  powder  and  carefully  washed  free  from  gritty  matter. 
It  is  then  mixed  with  the  finest  washed  clay,  and  the  pasty  mass  is 
forced  by  hydraulic  pressure  through  perforated  plates.  The  name 
"  graphite,"  from  the  Greek  to  write,  is  given  to  this  substance  on 
account  of  its  use  for  this  purpose.  It  was  formerly  supposed 
that  this  material  contained  lead,  hence  the  names  black-lead  and 
plumbago. 

When  powdered  graphite  is  subjected  to  prolonged  treatment  with  boiling 
nitric  acid  and  potassium  chlorate  it  undergoes  partial  oxidation,  and  is  con- 
verted into  a  greyish  crystalline  substance  which  was  termed  by  its  discoverer 
(Brodie)  graphitic  acid.  It  contains  carbon,  hydrogen,  and  oxygen,  and  is 


Carbon  289 

^» 

believed  to  have  a  composition  represented  by  the  formula  H4CnO5.  When 
heated,  this  compound  undergoes  a  very  curious  transformation.  If  a  frag- 
ment about  the  size  of  a  pea  is  heated  in  the  bottom  of  a  test-tube,  feeble  signs 
of  visible  combustion  are  seen,  and  a  light,  porous  black  mass  is  produced 
which  fills  and  overflows  the  tube.  Tnis  porous  mass  appears  to  be  pure 
graphite.  At  the  same  time  a  little  moisture  condenses  upon  the  tube. 

Graphite  is  largely  employed,  on  account  of  its  refractoriness, 
for  the  manufacture  of  the  so-called  plumbago  crucibles,  which 
consist  of  fireclay  mixed  with  finely-ground  graphite. 

Other  uses  to  which  graphite  is  put  are  for  glazing  or  polishing 
gunpowder,  especially  the  larger  grained  varieties  ;  as  a  lubricant 
for  machinery,  where  oil  is  inadmissible  on  account  of  high  tem- 
perature ;  for  electrotyping  processes,  and  also  as  a  coating  for 
ironwork,  to  prevent  rusting. 

AMORPHOUS  CARBON. 

This  non-crystalline  form  of  carbon  may  be  obtained  by  the 
decomposition  of  a  great  variety  of  carbon  compounds,  by  the 
process  known  as  destructive  distillation.  The  carbon  so  obtained 
differs  very  much  as  regards  its  purity,  according  to  the  particular 
organic  compound  used  for  its  preparation.  The  commonest  forms 
of  amorphous  carbon  to  be  met  with  are  lampblack  or  soot,  gas 
carbon,  coke,  charcoal,  animal  charcoal  or  bone  -  black.  None 
of  these  substances  is  pure  carbon  ;  animal  charcoal,  for  example, 
usually  containing  only  about  10  per  cent,  of  carbon. 

Lampblack. — This  substance  is  manufactured  by  burning  sub- 
stances rich  in  carbon,  and  which  burn  with  a  smoky  flame  (as 
turpentine,  petroleum,  or  tar),  with  a  limited  supply  of  air.  The 
smoke  is  passed  into  chambers  in  which  are  suspended  coarse 
blankets,  upon  which  the  soot  collects.  The  lampblack  always 
contains  hydrogen  in  the  form  of  hydrocarbons.  If  the  soot  be 
heated  to  redness  in  a  stream  of  chlorine,  this  hydrogen  can  be 
removed,  and  pure  amorphous  carbon  will  be  left. 

Lampblack  is  used  for  printers'  ink  and  for  black  paint. 

Gas  Carbon. — This  form  of  carbon  is  obtained  by  the  destruc- 
tive distillation  of  coal  in  the  manufacture  of  illuminating  gas. 
It  remains  in  the  retorts  as  an  extremely  hard  deposit,  lining  the 
roof  and  sides.  It  is  a  very  pure  carbon,  coming  second  to  puri- 
fied lampblack.  Its  specific  gravity  is  about  2.35.  Gas  carbon  is 
a  good  conductor  of  electricity,  and  is  extensively  used  for  the 
manufacture  of  carbon  rods  for  the  arc  light. 

T 


290  Inorganic  Chemistry 

Coke. — This  substance  differs  from  gas  carbon,  although  it  also 
is  obtained  in  the  process  of  coal  distilling.  It  contains  all  the 
inorganic  matter  which  constitutes  the  ash  of  the  coal,  and  also 
small  quantities  of  hydrogen,  nitrogen,  and  oxygen.  The  average 
amount  of  carbon  in  coke  is  about  91  per  cent. 

Charcoal. — The  purest  form  of  charcoal  is  obtained  by  the 
carbonisation  of  pure  white  sugar  and  the  subsequent  ignition 
of  the  charcoal  in  a  stream  of  chlorine  gas.  Charcoal  so  ob- 
tained has  a  specific  gravity  of  1.57.  Charcoal  in  a  much  less 
pure  condition  is  manufactured  from  wood.  The  methods 
by  which  the  carbonisation  of  wood  is  carried  out  are,  broadly, 
of  two  kinds  :  first,  those  in  which  access  of  air  is  permitted 
to  the  burning  material  ;  and,  second,  those  in  which  air  is  ex- 
cluded. 

The  first  of  these,  and  the  most  ancient,  is  generally  carried  on 
in  more  primitive  parts,  where  wood  is  plentiful.  The  wood  is 
piled  into  mounds  or  stacks,  which  are  built  with  some  care. 
They  are  set  on  fire  in  the  interior  by  means  of  a  lighted  bundle 
of  brushwood,  which  is  introduced  through  a  vertical  opening  or 
chimney,  left  for  this  purpose  in  the  centre  of  the  mound  during  its 
construction.  The  outside  of  the  heap  is  covered  with  brushwood, 
and  finally  with  turf,  in  order  to  regulate  the  access  of  air  to  the 
interior,  and  therefore  to  control  the  rate  of  combustion  of  the  wood. 
The  object  of  the  charcoal-burner  is  to  carbonise  the  wood  as 
slowly  as  possible.  In  this  process  there  is  great  liability  to  loss, 
by  the  too  rapid  combustion  of  the  wood  ;  and,  in  addition,  it  pos- 
sesses the  disadvantage  that  the  secondary  products,  such  as  the 
pyroligneous  acid,  tar,  &c.,  are  entirely  lost. 

Various  modifications  have  been  introduced  into  the  method  of 
coaling wood,  as  the  process  is  termed,  with  a  view  to  collect  these 
products. 

In  the  second  general  process  of  carbonising  wood,  the  material 
is  heated  in  ovens  or  retorts  from  the  outside,  no  air  being  admitted 
to  the  wood.  The  operation  is  very  similar  to  that  employed  in 
the  destructive  distillation  of  coal,  in  the  manufacture  of  coal  gas. 
In  these  methods  all  the  volatile  and  condensable  products  are 
collected  ;  among  these  are  water,  pyroligneous  acid,  wood  spirit, 
acetone,  and  fatty  oils.  The  non-condensable  products  consist 
mainly  of  such  gases  as  hydrogen,  carbon  monoxide,  carbon  di- 
oxide, marsh  gas,  and  acetylene. 

Animal  Charcoal. — Bone-black  is  obtained  by  the  carbonisa- 


Charcoal  291 

tion  of  bones  in  iron  retorts.     This  variety  of  charcoal  is  the  least 
pure  of  all  the  ordinary  forms  of  amorphous  carbon. 

Bone  contains  only  about  30  per  cent,  of  organic  matter,  the 
other  70  per  cent,  consisting  chiefly  of  calcium  phosphate,  asso- 
ciated with  small  quantities  of  magnesium  phosphate  and  calcium 
carbonate.  It  will  be  obvious,  therefore,  that  as  the  carbon  is 
derived  from  the  organic  matter,  the  amount  of  it  in  carbonised 
bones  must  be  small.  The  average  composition  of  animal  char- 
coal is  found  to  be — 

Carbon        .         .         .  .         .         .     10.0 

Calcium  phosphate     .  88.0 

Other  saline  substances  2.0 


i  oo.o 


Although  containing  relatively  so  small  an  amount  of  carbon, 
animal  charcoal  possesses  many  of  the  valuable  properties  of 
charcoal  in  a  highly  marked  degree,  owing  to  the  fact  that  it  con- 
tains its  carbon  disseminated  throughout  an  extremely  porous  mass 
of  calcium  phosphate. 

Properties  Of  Charcoal. — Charcoal  varies  very  considerably  in 
its  properties,  depending  upon  the  particular  wood  from  which  it 
is  obtained,  and  the  method  by  which  it  is  prepared.  Thus,  char- 
coal obtained  at  300°  is  a  soft,  brownish-black,  very  friable  material, 
having  an  igniting  point  as  low  as  380°.  On  the  other  hand, 
charcoal  prepared  at  very  high  temperatures  is  black  and  com- 
paratively dense,  and  requires  to  be  strongly  heated  in  order  to 
ignite  it. 

Under  ordinary  circumstances,  charcoal  burns  in  air  without  the 
formation  of  a  flame,  or  the  production  of  smoke.  At  high  tem- 
peratures, however,  the  combustion  of  charcoal  is  seen  to  be 
attended  by  a  flame.  This  is  probably  accounted  for  by  the  fact, 
that  as  the  temperature  at  which  the  combustion  of  carbon  takes 
place  is  raised  above  700°,  the  amount  of  carbon  monoxide  which 
is  formed  increases,  and  the  carbon  dioxide  decreases.*"* 

When  charcoal  is  thrown  upon  water  it  floats,  on  account  of  the 
r.ir  which  is  enclosed  within  its  pores.  The  specific  gravity  of 
charcoal  when  thus  filled  with  air  varies  from  0.106  (charcoal  made 

*  Ernst,  Chcmisches  Repertorium,  vol.  xvii.  p.  2. 


292  Inorganic  Chemistry 

from  the  ash)  to  0.203  (charcoal  from  the  birch).  If  the  air  be 
withdrawn  from  charcoal  it  sinks  in  water,  the  average  specific 
gravity  of  charcoal  itself  being  1.5. 

Ordinary  charcoal  is  a  bad  conductor  of  electricity,  but  its  con- 
ductivity is  greatly  increased  by  strongly  heating  the  charcoal  in 
closed  vessels. 

Charcoal  has  the  power  of  absorbing  gases  and  vapours  to  a 
remarkable  extent :  this  power,  which  is  exhibited  to  a  dif- 
ferent degree  by  the  various  kinds  of  charcoal,  is  due  to  the 
porosity  of  the  material,  whereby  it  exposes  a  very  large  sur- 
face ;  and  it  belongs  to  a  class  of  phenomena  known  as  surface 
action. 

If  a  fragment  of  charcoal,  recently  strongly  heated  to  expel  the 
air  from  its  pores,  be  passed  up  into  a  cylinder  of  ammonia  gas 
standing  in  a  trough  of  mercury,  the  ammonia  will  be  gradually 
absorbed  by  the  charcoal,  and  the  mercury  will  ascend  in  the 
cylinder.  Saussure  found  that  recently  heated  beech-wood  char- 
coal was  capable  of  absorbing  ninety  times  its  own  volume  of 
ammonia  gas  ;  while  Hunter,  by  employing  charcoal  made  from 
cocoa-nut  shell,  found  that  171.7  volumes  of  ammonia  were  absorbed 
by  i  volume  of  charcoal.  The  results  of  both  of  these  expe  *• 
ments  show  that  those  gases  are  absorbed  in  the  largest  quantities 
which  are  the  most  readily  liquefiable.  The  gas  so  held  by  the 
charcoal  is  in  a  highly  condensed  condition  upon  the  surface  of 
the  porous  mass.  Probably  in  the  case  of  easily  liquefied  gases, 
such  as  ammonia,  sulphur  dioxide,  and  others,  the  gases  are  par- 
tially liquefied  upon  the  surface  of  the  charcoal.  In  this  condensed 
state  the  gas  is  more  chemically  active  than  under  ordinary  condi- 
tions, and  charcoal  is  therefore  able  to  induce  many  striking  com- 
binations to  take  place.  Thus,  if  charcoal  be  allowed  to  absorb 
chlorine,  and  dry  hydrogen  be  then  passed  over  it,  the  chlorine  is 
capable  of  combining  with  the  hydrogen  even  in  the  dark,  with  the 
formation  of  hydrochloric  acid.  This  chemical  activity  of  gases, 
when  absorbed  by  charcoal,  is  strikingly  exemplified  in  the  case 
of  sulphuretted  hydrogen.  If  a  quantity  of  powdered  charcoal, 
which  has  been  saturated  with  sulphuretted  hydrogen,  be  brought 
into  oxygen,  the  rapid  combination  of  the  two  gases  is  attended 
with  the  development  of  so  much  heat  that  the  charcoal  bursts 
into  active  combustion.  In  the  same  way  a  mixture  of  air,  with 
10  or  15  per  cent,  of  sulphuretted  hydrogen,  may  be  passed  rapidly! 
through  a  tube,  about  a  metre  in  length,  filled  with  charcoal, 


Coal  293 

without  a  trace  of  sulphuretted  hydrogen  escaping  at  the  end.* 
Owing  to  this  property  charcoal  is  largely  employed  to  absorb 
noxious  gases,  the  atmospheric  oxygen  which  is  condensed  in  the 
pores  of  the  charcoal  oxidising  these  offensive  and  injurious  com- 
pounds ;  thus  sewer  ventilators  are  often  trapped  with  a  layer  of 
charcoal,  which  effectually  arrests  all  bad-smelling  gases. 

Charcoal  also  has  the  power  of  absorbing  colouring  matters  from 
solution  :  thus,  if  water  which  has  been  tinted  with  an  organic 
colouring  matter  be  shaken  up  with  powdered  charcoal  and  filtered, 
the  solution  will  be  found  to  be  entirely  decolourised.  The  variety 
of  charcoal  which  possesses  this  property  in  the  highest  degree  is 
animal  charcoal,  or  bone-black,  and  this  substance  is  largely  em- 
ployed in  many  manufacturing  processes,  such  as  sugar-refining, 
tin  order  to  remove  all  colouring  matter  from  the  liquid. 

Charcoal  under  ordinary  conditions  is  unacted  upon  by  the  air, 
but  when  the  temperature  is  raised  it  enters  into  active  combus- 
tion, forming  carbon  dioxide.  In  an  extremely  divided  condition, 
however,  carbon  is  capable  of  combining  spontaneously  with  the 
oxygen  of  the  air,  and  with  so  much  energy  as  to  take  fire. 

Coal. — The  carbonaceous  minerals  that  are  included  under  the 
.  f  jame  coal  are  an  impure  form  of  carbon,  containing  compounds 
of  carbon  with  hydrogen  and  oxygen.  Coal  is  the  final  result  of  a 
series  of  decomposition  changes  which  have  been  undergone  by 
vegetable  matter  of  the  remote  past,  the  process  having  extended 
over  long  geological  periods.  During  this  prolonged  process  a  por- 
tion of  the  carbon  and  hydrogen  is  eliminated  as  marsh  gas,  and 
large  quantities  of  this  gas  are  found  associated  with,  and  occluded 
in,  coal. 

Broadly  speaking,  the  numerous  varieties  of  coal  may  be  divided 
into  soft  or  bituminous,  and  hard  or  anthradtic. 

The  former  are  employed  for  the  manufacture  of  coal  gas  and 
for  ordinary  domestic  purposes  ;  they  burn  with  a  smoky  flame,  and 
evolve  large  quantities  of  gases  and  volatile  vapours  on  combus- 
tion or  distillation.  Anthracite  coal  is  much  harder,  ignites  with 
more  difficulty,  and  burns  with  the  formation  of  very  little  flame  or 
smoke.  It  contains  a  higher  percentage  of  carbon,  and  gives  out 
great  heat  on  combustion,  and  is  employed  largely  as  a  steam- 
coal. 

*  "  Chemical  Lecture  Experiments,"  394-396,  new  ed. 


294 


Inorganic  Chemistry 


The  following  table  shows  the  average  composition  of  coals  from 
various  sources,  and  the  general  difference  between  coals  of  the 
two  main  classes  : — 


c 

d 

D 

he 

g 

g. 

3 

j^ 

• 

.   Locality. 

J 

° 

R 

& 

J3 

4 

CC 

^ 

0 

pq 

O 

* 

M 

^Northumberland 

3    S    1 

81.41 

5-83 

7.90 

2.05 

0.74 

2.07 

».3S 

66.70 

*j  o  j  \Vpipc 
ffl.S|wales 

83.78 

4-79 

4-15 

0.98 

1-43 

4.91 

72.6O 

s 

^Staffordshire 

78.57 

5-29 

12.88 

1.84 

o-39 

1.03 

11.29 

57-21 

d       TS  Wales 

92  56 

•a  qo 

2   I?'? 

i  =;8 

£  a;  I  °* 

JE;  °  (Pennsylvania 

QO  4^ 

2  4^ 

2  4? 

4.6 

CHAPTER  IX 

CARBON  COMPOUNDS 

THE  compounds  of  the  element  carbon  are  so  numerous  that  it 
has  been  found  convenient  to  constitute  the  study  of  these  sub- 
stances a  separate  branch  of  chemistry.  In  the  early  history  of 
the  science  it  was  believed  that  there  were  a  large  number  of 
substances  which  could  only  be  obtained  as  the  product  of  living 
organisms.  They  were  known  to  be  elaborated  by  the  action  of 
life,  or,  as  it  was  termed,  the  vital  force,  and  it  was  believed  that 
owing  to  some  inherent  specific  quality  belonging  to  this  vital  force 
the  substances  produced  by  its  action  were  distinct  from  such 
substances  as  could  be  prepared  by  any  laboratory  processes.  To 
denote  this  distinction,  the  term  organic  was  applied  to  those  things 
which  were  known  to  be  the  products  of  living  organisms,  and 
other  compounds  were  distinguished  as  inorganic  substances.  This 
distinction  received  its  deathblow  in  1828,  when  Wohler  produced, 
by  purely  laboratory  processes,  one  of  the  most  typical  of  all  organic 
compounds,  namely,  urea.  The  names  "  organic  "  and  "  inorganic " 
chemistry  are  still  retained,  but  their  old  significance  is  entirely 
gone,  as  no  distinction  is  to-day  recognised  between  products  elabo- 
rated by  the  action  of  life  and  those  which  can  be  synthetically 
produced. 

Speaking  broadly,  organic  chemistry  may  be  described  as  the 
chemistry  of  the  carbon  compounds.  Nevertheless,  although  it  is 
quite  true  that  all  "  organic  compounds  "  contain  carbon,  it  has  not 
been  found  expedient  to  include  in  the  category  of  organic  com- 
pounds all  compounds  containing  carbon.  Not  because  there  is 
any  intrinsic  difference  in  these  compounds,  but  merely  from  con- 
siderations of  convenience.  The  following  may  be  mentioned  as 
examples  of  such  compounds  as  are  not  regarded  as  belonging  to 
the  "  organic "  division  :  compounds  of  carbon  with  the  metals, 
namely,  the  so-called  carbides,  of  which  cast  iron  and  calcium 
carbide  are  familiar  cases  ;  the  compounds  of  carbon  with  sulphur 


2g6  Inorganic  Chemistry 

and  the  extensive  series  of  thio-carbonates  ;  carbon  monoxide 
and  the  compounds  formed  by  its  direct  union  with  non-metals 
(e.g.  carbonyl  chloride,  &c.)  and  with  metals  (e.g.  nickel  carbonyl, 
&c.)  ;  and  lastly,  carbon  dioxide  and  all  the  multitude  of  metallic 
carbonates.  Obviously,  therefore,  the  broad  distinction  above 
mentioned  must  not  be  regarded  as  a  definition.  Indeed,  it 
may  be  said  that  no  exact  definition  of  an  "  organic"  compound  has 
ever  been  framed,  and  we  have  to  accept  the  general  statement 
that  "  organic "  chemistry  is  the  chemistry  of  the  carbon  com- 
pounds with  certain  generally  acknowledged  exceptions. 

Amongst  the  compounds  of  carbon  which  will  be  briefly  treated 
of  in  these  chapters,  there  will  be  included  three  which  all  chemists 
agree  to  regard  as  organic  substances  :  these  are  methane  (marsh 
gas),  CH4  ;  ethylene,  C2H4  ;  and  acetylene,  C2H2.  These  three  com- 
pounds play  an  important  part  in  our  ordinary  illuminating  flames 
and  in  coal  gas. 

COMPOUNDS  OF  CARBON  WITH  OXYGEN. 

There  are  two  well-known  oxides  of  carbon,  both  of  which  are 
colourless  gases,  viz.  : — 

Carbon  monoxide CO, 

Carbon  dioxide CO2. 

CARBON  MONOXIDE. 

Formula,  CO.     Molecular  weight=28.     Density =14. 

Modes  of  Formation.— ( i.)  Carbon  monoxide  is  formed  when 
carbon  dioxide  is  passed  over  charcoal  heated  to  bright  redness — 

C02  +  C  =  2CO. 

The  same  result  is  obtained  when  a  slow  stream  of  air  or  oxygen 
is  passed  over  red-hot  charcoal  contained  in  a  tube.  The  first 
action  of  the  air  on  coming  in  contact  with  the  carbon  is  to  form 
carbon  dioxide,  which,  passing  over  the  remainder  of  the  heated 
material,  is  deprived  of  a  portion  of  its  oxygen  according  to  the 
above  equation.  This  operation  goes  on  in  an  ordinary  fire-grate  : 
the  air,  on  first  gaining  access  to  the  burning  coal  or  coke,  causes 
the  complete  oxidation  of  a  portion  of  the  carbon  to  carbon  dioxide  ; 
and  as  this  gas  passes  through  the  mass  of  red-hot  carbon  it  is 
reduced  to  the  lower  oxide,  which  either  escapes  with  the  other 


Carbon  Monoxide  297 

products  of  combustion  or  becomes  ignited  and  burns  with  a 
lambent  bluish  flame  such  as  may  frequently  be  noticed  upon  the 
top  of  a  "  clear  "  fire. 

(2.)  When  steam  is  passed  over  strongly  heated  carbon  a  mixture 
of  carbon  monoxide  and  hydrogen  is  produced.  This  mixture, 
known  as  water  gas,  is  employed  in  many  manufacturing  processes 
as  a  gaseous  fuel  — 


(3.)  Carbon  monoxide  is  also  formed  by  the  action  of  carbon 
dioxide  upon  red-hot  iron  — 


=  Fe3O4+4CO. 

(4.)  Or  by  strongly  heating  either  carbon  or  iron  with  a  car- 
bonate, such  as  calcium  carbonate,  which  is  capable  of  yielding 
carbon  dioxide,  thus  — 


4CaCO3  +  3Fe  =  Fe3O4  +  4CaO  +  4CO. 

(5.)  Carbon  monoxide  is  most  conveniently  prepared,  by  the 
decomposition  of  certain  organic  compounds  by  means  of  sulphuric 
acid.  Thus,  when  formic  acid,  or  a  formate,  is  acted  upon  by  sul- 
phuric acid,  the  sulphuric  acid  withdraws  the  elements  of  water 
from  the  molecule  of  formic  acid,  and  leaves  carbon  monoxide  — 

H'COOH-H20  =  CO. 

(6.)  By  a  similar  decomposition,  oxalic  acid  yields  a  mixture  of 
carbon  monoxide  and  carbon  dioxide  in  equal  volumes  — 


C2H2O4-H20  = 

The  carbon  dioxide  is  readily  removed  from  the  mixture,  by  passing 
the  gases  through  a  solution  of  sodium  hydroxide  (caustic  soda), 
in  which  carbon  dioxide  is  absorbed  with  the  formation  of  sodium 
carbonate. 

(7.)  The  method  usually  employed  when  carbon  monoxide  is 
required  for  experimental  purposes  consists  in  heating  a  mixture 
of  one  part  by  weight  of  crystallised  potassium  ferrocyanide  (yellow 
prussiate  of  potash)  with  ten  parts  of  strong  sulphuric  acid  in  a 
capacious  flask,  when  the  following  reaction  takes  place  — 

K4FeC6N6  +  6H2SO4  +  6H2O  =  2K2SO4  +  FeSO4 
+  3(NH4)2SO4  ~ 


298  Inorganic  Chemistry 

The  six  molecules  of  water  required  by  the  reaction  are  derived 
partly  from  the  acid  employed  and  partly  from  the  salt,  which 
contains  three  molecules  of  water  of  crystallisation.* 

Properties.  —  Carbon  monoxide  is  a  colourless,  tasteless  gas, 
having  a  faint  smell.  It  is  only  slightly  soluble  in  water,  its  co- 
efficient of  absorption  at  o°  being  0.03287.  It  burns  in  the  air  with 
a  characteristic  pale-blue  flame,  forming  carbon  dioxide  — 


When  mixed  with  half  its  own  volume  of  oxygen,  and  inflamed, 
the  mixture  explodes  with  some  violence.t  If  the  two  gases  be 
confined  in  a  eudiometer  standing  over  mercury,  and  be  rendered 
absolutely  free  from  aqueous  vapour  by  powerful  desiccating  agents, 
no  explosion  will  take  place  upon  the  passage  of  an  electric  spark 
through  the  mixture.  And  in  the  same  way,  if  carbon  monoxide, 
which  has  been  deprived  of  all  aqueous  vapour,  be  burned  from  a 
jet  in  the  air,  and  the  jet  be  lowered  into  a  cylinder  containing  air 
which  has  been  similarly  dried,  the  flame  will  be  extinguished 
(see  page  191). 

Carbon  monoxide  is  an  extremely  poisonous  gas  :  very  small 
quantities  present  in  the  air  rapidly  give  rise  to  headache  and 
giddiness,  and  if  inhaled  for  a  length  of  time,  or  if  taken  into  the 
lungs  in  a  less  dilute  condition,  insensibility  and  death  quickly 
follow.  The  deaths  that  have  resulted  from  the  use  of  unventi- 
lated  fires  —  either  of  charcoal  or  coke,  or  in  some  cases  of  coal  gas 
—  in  dwelling-rooms,  have  been  due  to  the  escape  of  this  poisonous 
gas  into  the  air.  The  extremely  deadly  nature  of  the  after-damp 
resulting  from  a  colliery  explosion  is  due  to  the  presence  of  carbon 
monoxide  in  the  carbon  dioxide  which  is  formed  as  a  product  of 
the  combustion. 

The  poisonous  action  of  this  gas  is  due  to  its  absorption  by  the  blood,  with 
the  formation  of  a  bright  red  compound,  to  which  the  name  carboxy-hcemo- 
globin  is  applied.  Blood  so  charged  appears  to  be  unable  to  fulfil  its  function 
of  absorbing  and  distributing  oxygen  throughout  the  system.  This  carboxy- 
hsemoglobin  gives  a  characteristic  absorption  spectrum,  which  furnishes  a 
ready  method  of  detection  in  cases  of  poisoning  from  the  inhalation  of  carbon 
monoxide. 


*  "  Chemical  Lecture  Experiments,"  new  ed.,  435-439. 

t  The  rate  at  which  the  combustion  is  propagated  throughout  a  mixture  of 
carbon  monoxide  and  oxygen  is  much  slower  than  through  hydrogen  and 
oxygen.  Bunsen  has  estimated  it  at  less  than  i  metre  per  second. 


Carbon  Monoxide  299 

Carbon  monoxide  is  one  of  the  most  difficultly  liquefiable  gases, 
its  critical  temperature  being  -  136°. 

At  high  temperatures  this  gas  is  a  powerful  reducing  agent, 
uniting  with  another  atom  of  oxygen  to  form  carbon  dioxide. 
This  fact  is  made  use  of  in  many  metallurgical  processes  for  re- 
ducing the  oxides  of  the  metals  to  the  metallic  state. 

Carbon  monoxide  is  absorbed  at  ordinary  temperatures  by  a 
solution  of  cuprous  chloride,  forming  the  compound  COCu2Cl2. 

At  a  temperature  of  boiling  water,  carbon  monoxide  is  slowly 
absorbed  by  solid  potassium  hydroxide,  with  the  formation  of 
potassium  formate  — 

H-COOK. 


Carbon  monoxide  unites  directly  with  chlorine,  under  the  in- 
fluence of  sunlight,  forming  the  compound  known  as  phosgene  gas^ 
or  carbonyl  chloride  — 

CO  +  C12=COC12. 

If  the  two  gases  are  mixed  in  equal  volumes,  and  kept  in  the 
dark,  no  action  takes  place,  but  on  exposure  to  sunlight  they  com- 
bine, and  the  yellowish  colour  due  to  the  chlorine  will  disappear. 
On  opening  the  vessel  in  moist  air,  clouds  of  hydrochloric  acid 
are  formed,  owing  to  the  decomposition  of  carbonyl  chloride  by 
the  moisture,  according  to  the  equation  — 


Carbonyl  chloride  may  be  readily  condensed  to  a  liquid,  its 
boiling-point  being  +  8°. 

Carbon  monoxide  unites  directly  with  certain  metals,  giving  rise 
to  compounds  which  possess  some  very  remarkable  properties, 
and  to  which  the  name  metallic  carbonyls  has  been  applied  by 
their  discoverer.* 

When  carbon  monoxide  is  allowed  to  stream  slowly  over  metallic 
nickel  (obtained  by  the  reduction  of  nickel  oxide  in  a  stream  of 
hydrogen),  the  gas  is  absorbed  by  the  finely-divided  metal,  forming 
a  compound  having  the  composition  Ni(CO)4.  If  the  issuing  gas 
be  passed  through  a  cooled  tube.,  the  nickel  carbonyl  condenses 
to  a  colourless,  mobile,  highly  refracting  liquid,  having  a  specific 
gravity  at  o°  of  1.356,  and  boiling  at  43°  under  a  pressure  of 
751  mm.t 

*  Mond,  1890. 
f  See  "Chemical  Lecture  Experiments,"  new  ed.,  446-448. 


300  Inorganic  Chemistry 

Nickel  carbonyl  vapour  burns  with  a  luminous  flame,  which 
produces  a  black  deposit  of  metallic  nickel  when  a  cold  porcelain 
dish  is  depressed  upon  the  flame.  The  gas  is  decomposed  into 
nickel  and  carbon  monoxide  if  passed  through  a  hot  glass  tube, 
the  nickel  being  deposited  as  a  bright  metallic  mirror  upon  the 
glass— 

Ni(CO)4  = 


A  similar  compound  of  carbon  monoxide  and  iron  has  also  been  obtained, 
having  the  composition  Fe(CO)5.  Iron  carbonyl  is  a  pale-yellow,  viscous 
liquid,  boiling  at  102.  8°  under  a  pressure  of  749  mm.  Its  specific  gravity  at 
18°  is  1.4664.  When  heated  to  180°  the  vapour  is  decomposed,  iron  being 
deposited  and  carbon  monoxide  being  evolved.  This  compound  has  been  found 
in  iron  cylinders  in  which  the  so-called  -water  gas  (a  mixture  of  H  and  CO)  has 
been  stored  under  pressure  for  a  length  of  time  ;  it  is  also  said  to  be  present  in 
minute  quantities  in  coal  gas. 


CARBON  DIOXIDE. 

Formula,  CO2.     Molecular  weight =44.     Density =22. 

History. — Van  Helmont  in  the  seventeenth  century  was  the  first 
to  distinguish  between  this  gas  and  ordinary  air  :  he  observed  that  it 
was  formed  during  the  processes  of  combustion  and  fermentation, 
and  he  applied  to  it  the  name  gas  sylvestre.  Black  showed  that 
this  gas  was  a  constituent  of  what  in  his  day  were  known  as  the 
mild  alkalis  (alkaline  carbonates),  and  on  account  of  its  being  so 
combined,  or  fixed,  in  these  substances,  he  named  the  gas  fixed 
air.  Lavoisier  first  proved  its  true  chemical  composition  to  be 
that  of  an  oxide  of  carbon. 

Occurrence. — Carbon  dioxide  is  a  constant  constituent  of  the 
atmosphere,  being  present  to  the  extent  of  about  3  volumes  in 
10,000  volumes  of  air.  It  is  also  found  in  solution  in  all  spring- 
water,  which  is  sometimes  so  highly  charged  with  this  gas  under 
pressure  that  the  water  is  effervescent,  or  "  sparkling,"  from  the 
escape  of  the  gas.  Carbon  dioxide  is  evolved  in  large  quantities 
from  vents  and  fissures  in  the  earth  in  volcanic  districts.  The 
well-known  Poison  Valley  in  Java,  which  is  an  old  volcanic  crater, 
and  the  Grotto  del  Cane  near  Naples,  owe  their  peculiar  pro- 
perties to  the  discharge  into  them  of  large  quantities  of  carbon 
dioxide  from  such  subterranean  sources. 


Carbon  Dioxide 


301 


Modes  Of  Formation.—  (  i.)  Carbon  dioxide*  is  produced  when 
carbon  is  burnt  with  a  free  supply  of  air  or  oxygen  — 

C  +  O2  =  CO2. 

If  an  insufficient  supply  of  oxygen  be  employed,  carbon  mon- 
oxide is  produced  at  the  same  time. 

(2.)  When  limestone  or  chalk  is  strongly  heated,  as  in  the 
process  of  burning  lime,  carbon  dioxide  is  evolved  in  large 
quantities  — 

CaCO3  = 


(3.)  In  the  ordinary  processes  of  fermentation,  and  during  the 
decay  of  many  organic  substances,  carbon  dioxide  is  also  formed. 


FIG.  62. 

Thus,  when  sugar  undergoes  alcoholic  fermentation  by  means  of 
yeast,  the  sugar  is  converted  into  alcohol  and  carbon  dioxide — 

C12H22On  +  H20  =  4C2H6O  +  4CO2. 

(4.)  Carbon  dioxide  is  formed  during  the  process  of  respiration  ; 
also  by  the  combustion  of  all  ordinary  fuels,  and  of  any  compound 
containing  carbon,  such  as  candles,  oils,  gas,  &c. 

(5.)  For  experimental  purposes,  carbon  dioxide  is  most  readily 
obtained  by  the  decomposition  of  a  carbonate  by  means  of  a 
stronger  acid.  The  effervescence  that  results  from  the  action  of 
tartaric  acid  upon  sodium  bicarbonate,  in  an  ordinary  Seidlitz 

*  Experiments  on  carbon  dioxide,  Nos.  400-434,  "  Chemical  Lecture  Ex- 
periments," new  ed. 


302 


Inorganic  Chemistry 


powder,  is  due  to  the  disengagement  of  this  gas.  The  most  con- 
venient carbonate  for  the  preparation  of  this  gas  is  calcium 
carbonate,  in  one  of  its  many  naturally  occurring  forms,  such  as 
marble,  limestone,  or  chalk.  Fragments  of  marble  are  for  this 
purpose  placed  in  a  two-necked  bottle  (Fig.  62),  with  a  quantity  of 
water,  and  strong  hydrochloric  acid  is  added  by  means  of  the 
funnel-tube.  A  rapid  effervescence  takes  place  owing  to  the 


FIG.  63. 


FIG.  64. 


elimination  of  the  gas,  and  a  solution  of  calcium  chloride  remains 
in  the  bottle  — 


If  sulphuric  acid  be  substituted  for  hydrochloric  acid,  the  frag- 
ments of  marble  rapidly  become  coated  with  a  crust  of  insoluble 
calcium  sulphate,  which  soon  prevents  the  further  action  of  the 
acid,  and  therefore  puts  an  end  to  the  reaction  :  by  employing 
finely  powdered  chalk,  however,  instead  of  lumps  of  calcium  car- 
bonate, this  difficulty  is  obviated.  This  gas  is  largely  manu- 
factured from  these  materials. 

Properties.  —  Carbon  dioxide  is  a  colourless  gas,  having  a  feeble 


Carbon  Dioxide  303 

acid  taste  and  a  faint  and  pleasantly  pungent  smell.  It  is  incap- 
able of  supporting  either  combustion  or  respiration;  a  burning 
taper  is  instantly  extinguished,  and  an  animal  speedily  dies  when 
introduced  into  this  gas.  Although  carbon  dioxide  is  not  such  a 
poisonous  compound  as  the  monoxide,  it  nevertheless  does  exert 
a  direct  poisonous  effect  upon  the  system,  and  death  caused  by 
this  gas  is  not  merely  due  to  the  absence  of  oxygen.  The  pro- 
longed inhalation  of  air  containing  only  a  very  slightly  increased 
amount  of  carbon  dioxide  has  a  distinctly  lowering  effect  upon 
the  vitality. 

Carbon  dioxide  is  a  heavy  gas,  being  about  one  and  a  half 
times  heavier  than  air.  On  this  account  it  may  readily  be  col- 
lected by  displacement.  By  virtue  of  its  great  density  it  may  be 
poured  from  one  vessel  to  another,  much  in  the  same  way  as  an 
ordinary  liquid  :  thus,  if  a  large  bell-jar  be  filled  with  the  gas  by 
displacement,  a  beaker-full  may  be  drawn  up,  as  water  from  a 
well  (Fig.  63).  If  the  gas  so  drawn  up  be  poured  into  a  similar 
beaker,  suspended  from  the  arm  of  a  balance,  and  counterpoised, 
the  weight  of  the  gas  will  be  evident  by  the  disturbance  of  the 
equilibrium  of  the  system. 

If  a  soap  bubble  be  allowed  to  fall  into  a  large  jar  filled  with 
carbon  dioxide,  it  will  be  seen  to  float  upon  the  surface  of  the 
dense  gas  (Fig.  64).  The  power  of  carbon  dioxide  to  extinguish 
flame  is  so  great,  that  a  taper  will  not  burn  in  air  in  which  this  gas 
is  present  to  the  extent  of  2.5  per  cent.,  and  in  which  the  oxygen 
is  reduced  to  18.5  per  cent.  For  this  reason  a  comparatively  small 
quantity  of  carbon  dioxide,  brought  into  the  air  surrounding  a  burn- 
ing body,  is  capable  of  extinguishing  the  flame.  This  property  has 
been  put  to  valuable  service  in  the  construction  of  numerous  con- 
trivances for  extinguishing  fire,  such  as  the  "  extincteur."  This  is 
a  metal  vessel  containing  carbon  dioxide  under  pressure,  the  gas 
having  been  generated  within  the  closed  apparatus  by  the  action 
of  dilute  sulphuric  acid  upon  sodium  carbonate.  A  stream  of  the 
gas,  projected  judiciously  upon  a  moderate  conflagration  in  a 
dwelling,  readily  extinguishes  the  fire.  This  property  may  be 
illustrated  by  inflaming  a  quantity  of  turpentine  in  a  dish,  and 
pouring  upon  the  flames  a  quantity  of  carbon  dioxide  contained 
in  a  large  bell-jar  (Fig.  65),  when  it  will  instantly  extinguish  the 
conflagration. 

Although  carbon  dioxide  is  incapable  of  supporting  combus- 
tion in  the  ordinary  sense,  certain  metals  are  capable  of  burn- 


304 


Inorganic  Chemistry 


ing  in  this  gas.  Thus,  a  fragment  of  potassium  when  heated 
in  this  gas  burns  brightly,  forming  potassium  carbonate  with 
the  deposition  of  carbon  — 

=  2KCO  +  C. 


When  carbon  dioxide  is  passed  into  a  solution  of  calcium 
hydroxide  (lime  water)  a  turbidity  at  once  results,  owing  to  the 
precipitation  of  insoluble  calcium  carbonate  or  chalk  — 


This  reaction  furnishes  the  readiest  means  for  the  detection  of 


FIG.  65. 

carbon  dioxide.  Thus,  if  the  gas  obtained  by  any  of  the  modes  of 
formation  described  be  passed  into  clear  lime  water,  the  formation 
of  this  white  precipitate  of  chalk  is  proof  that  the  gas  is  carbon 
dioxide.  By  this  test  it  may  readily  be  shown  that  carbon  dioxide 
is  a  product  of  respiration,  by  merely  causing  the  exhaled  breath 
to  bubble  through  a  quantity  of  lime  water,  which  will  quickly  be 
rendered  turbid. 

Carbon  dioxide  is  moderately  soluble  in  water.  At  the  ordinary 
temperature  water  dissolves  about  its  own  volume  of  this  gas. 

The  coefficient  of  absorption  at  o°  is  1.7967,  the  solubility  de- 
creasing with  rise  of  temperature  in  accordance  with  the  inter- 
polation formula — 

c=  1.7967  -  0.07761^+  o.oo  1 6424/2. 


Carbon  Dioxide  305 

Carbon  dioxide  shows  a  slight  departure  from  Henry's  law 
(see  page  143),  when  the  pressures  are  greater  than  that  of  the 
atmosphere.  Thus,  when  the  pressure  is  doubled,  the  amount  dis- 
solved is  slightly  more  than  doubled.  The  solubility  of  carbon 
dioxide  in  water,  and  its  increased  solubility  under  pressure,  is 
illustrated  in  the  ordinary  aerated  waters.  Water  under  a  pres- 
sure of  several  atmospheres  is  saturated  with  the  gas,  and  upon 
the  release  of  this  pressure  by  the  withdrawal  of  the  cork  the 
excess  of  gas,  over  and  above  that  which  the  water  can  dissolve  at 
the  ordinary  pressure,  escapes  with  the  familiar  effervescence.  In 
a  similar  manner  the  natural  aerated  waters  have  thus  become 
charged  with  carbon  dioxide,  under  subterranean  pressure,  and 
when  such  waters  come  to  the  surface  the  dissolved  gas  begins 
to  make  its  escape. 

The  solution  of  carbon  dioxide  in  water  is  feebly  acid,  turning 
blue  litmus  to  a  port-wine  red  colour,  characteristically  different 
from  the  scarlet  red  given  by  stronger  acid:?.  This  acid  may  be 
regarded  as  the  true  carbonic  acid — 

C02  +  H20  =  H2C03. 

A  recently-made  sample  of  aerated  water  is  seen  to  effervesce 
more  briskly  and  give  off  the  dissolved  gas  more  rapidly  than 
specimens  that  have  been  long  preserved.  In  process  of  time  the 
dissolved  carbon  dioxide  gradually  combines  with  the  water,  with 
the  formation  of  carbonic  acid,  an  unstable  compound  which  slowly 
decomposes  into  carbon  dioxide  and  water,  especially  at  a  slight 
elevation  of  temperature.  Many  of  the  naturally  occurring  aerated 
waters,  such  as  Apollinaris,  when  opened  exhibit  scarcely  any 
effervescence,  but  give  off  carbon  dioxide  gradually.  Such  waters 
have  in  all  probability  been  exposed  to  pressure  for  a  great  length 
of  time,  and  their  dissolved  carbon  dioxide  has  almost  entirely 
combined  to  form  carbonic  acid.  When  such  a  water  is  gently 
warmed  a  rapid  stream  of  gas  is  evolved. 

When  carbon  dioxide  is  strongly  heated,  as  by  the  passage  of 
electric  sparks,  it  is  partially  dissociated  into  carbon  monoxide 
and  oxygen.  This  decomposition  is  never  complete  ;  for  when 
the  amount  of  these  two  gases  in  the  mixture  reaches  a  certain 
proportion,  they  reunite  to  form  carbon  dioxide,  and  a  point  of 
equilibrium  is  reached  when  as  many  molecules  are  united  as  are 
dissociated  in  the  same  time. 

Liquid  Carbon  Dioxide.— Carbon  dioxide  is  easily  liquefied. 

u 


306  Inorganic  Chemistry 

At  -5°  it  requires  a  pressure  of  30.8  atmospheres  ;  at  +5°,  40.4 
atmospheres;  while  at  +15°  a  pressure  of  52.1  atmospheres  is 
required. 

Faraday  first  liquefied  this  gas,  by  introducing  into  a  strong  bent 
glass  tube  a  quantity  of  sulphuric  acid  and  a  few  lumps  of  ammo- 
nium carbonate,  which  were  prevented  from  touching  the  acid  by 
means  of  a  plug  of  platinum  foil.  The  tube  was  then  hermetically 
sealed,  and  the  acid  allowed  gently  to  come  in  contact  with  the 
carbonate,  which  was  at  once  decomposed  with  the  formation  of 
ammonium  sulphate  and  carbon  dioxide.  By  the  internal  pres- 
sure exerted  by  the  evolved  gas,  aided  by  the  application  of  cold 
to  one  end  of  the  bent  tube,  the  gas  condensed  to  a  colourless 
liquid. 

Large  quantities  of  this  liquefied  gas  were  obtained  by  Thilorier 
by  a  precisely  similar  method,  the  experiment  being  performed  in 
strong  wrought-iron  vessels. 

Liquid  carbon  dioxide  is  to-day  manufactured  on  a  large  scale, 
by  pumping  the  gas  into  steel  cylinders  by  means  of  powerful 
compression  pumps.  The  enormous  volumes  of  carbon  dioxide 
evolved  in  the  process  of  brewing,  and  which  until  quite  recently 
were  allowed  to  escape  into  the  atmosphere,  are  now  utilised  for 
this  purpose.  The  gas,  as  it  is  evolved  from  the  fermenting  vats, 
is  washed  and  purified,  and  pumped  into  steel  bottles  for  the 
market.  In  this  form  the  gas  is  largely  employed  by  manufac- 
turers of  aerated  waters,  and  also  as  the  refrigerating  agent  in 
"  cold  storage." 

Liquid  carbon  dioxide  is  a  colourless  and  extremely  mobile 
liquid,  which  floats  upon  water  without  mixing.  It  boils  at  -80° 
under  atmospheric  pressure. 

When  heated,  liquid  carbon  dioxide  expands  at  a  more  rapid 
rate  than  a  gas,  its  coefficient  of  expansion  being  greater  than  that 
of  any  known  substance.  Its  rapid  change  of  volume  is  seen  by 
the  following  figures  : — 

95  volumes  at  -  10°  become 
100       „         „         o°       „ 
1 06       „         „    +10°       „ 

114          „  „     +20° 

The  critical  temperature  of  carbon  dioxide  is  31.35°.  If  the  liquid 
be  heated  to  this  point,  it  passes  into  the  gaseous  state  without  any 
change  of  volume.  The  line  of  demarcation  between  the  liquid 


Carbon  Dioxide 


307 


and  gas  in  the  tube  gradually  fades  away,  and  the  tube  appears 
filled  with  gas.  Above  this  temperature  no  additional  pressure  is 
able  to  liquefy  the  gas.  On  once  more  cooling  the  tube,  when  the 
critical  point  is  passed  the  liquid  again  appears,  and  the  dividing 
line  between  it  and  the  gas  is  once  more  sharply  defined. 

Solid  Carbon  Dioxide.— When  liquid  carbon  dioxide  is  allowed 
to  escape  into  the  air  the  absorption  of  heat  due  to  its  rapid  eva- 
poration causes  a  portion  of  the  liquid  to  solidify.  This  solid  is 
most  conveniently  collected  by  allowing  the  jet  of  liquid  to  stream 
into  a  round  metal  box  (Fig.  66),  in  which  it  is  caused  to  revolve  by 
being  made  to  impinge  upon  the  curved  tongue  of  metal.  The  box 
is  furnished  with  hollow  wooden  handles,  through  which  the  gas 
makes  its  escape.  Considerable  quantities  of  the  frozen  carbon 
dioxide  can  in  this  way  be  collected  in  a  few  minutes. 

On  a  larger  scale  the  brass  box  is 
substituted  by  a  canvas  bag,  which  is 
simply  tied  over  the  nozzle  of  the  cylinder 
containing  the  liquefied  gas,  and  a  rapid 
stream  of  the  liquid  allowed  to  escape 
into  it. 

Solid  carbon  dioxide  is  a  soft,  white, 
snow-like  substance.  When  exposed  to 
the  air  it  quickly  passes  into  gas,  without 
going  through  the  intermediate  state  of 
liquidity. 

Solid  carbon  dioxide  is  readily  soluble 
in  ether,  and  this  solution  constitutes  one 
of  the  most  convenient  sources  of  cold. 
A  large  number  of  gases  can  readily  be 
liquefied  by  being  passed  through  tubes 
immersed  in  this  freezing-mixture.  When  this  ethereal  solution 
is  rapidly  evaporated  its  temperature  can  be  lowered  to  —  no". 

"  Carbonic  acid  snow,"  as  this  substance  is  sometimes  termed,  is 
now  an  article  of  commerce,  the  compound  being  sent  into  the 
market  in  this  form  to  avoid  the  cost  of  the  carriage  of  the  neces- 
sarily heavy  steel  bottles  containing  the  liquid. 

Composition  of  Carbon  Dioxide.— When  carbon  burns  in 
oxygen,  the  oxygen  undergoes  no  change  in  volume  in  being  trans- 
formed into  carbon  dioxide.  The  volume  of  carbon  dioxide  produced 
is  the  same  as  that  of  the  oxygen  which  is  required  for  its  produc- 
tion. This  may  be  shown  by  means  of  the  apparatus  (Fig.  67). 


FIG.  66. 


308 


Inorganic  Chemistry 


The  bulb  of  the  U-tube  is  filled  with  oxygen,  and  the  stopper, 
which  carries  a  small  bone-ash  crucible  upon  which  a  fragment 
of  charcoal  is  placed,  is  lowered  into  position.  The  charcoal  is 
ignited  by  means  of  a  thin  loop  of  platinum  wire,  as  shown  in  the 
figure,  which  can  be  heated  by  an  electric  current.  As  the  carbon 
burns  the  heat  causes  a  temporary  expansion  of  the  included  gas  ; 
but  after  the  combustion  is  complete  and  the  apparatus  has 
cooled,  the  level  of  mercury  will  be  found  to  be  undisturbed. 
Carbon  dioxide,  therefore,  contains  its  own  volume  of  oxygen. 
From  this  experiment  the  composition  of  carbon  dioxide  by  weight 
can  be  deduced.  One  litre  of  carbon  dioxide  weighs  22  criths  ; 
deducting  from  this  the  weight  of  I  litre  of  oxygen,  viz.,  16 
criths,  we  get  6  as  a  remainder.  Six  parts  by 
weight  of  carbon,  therefore,  combine  with  16 
parts  by  weight  of  oxygen  to  form  22  parts 
of  carbon  dioxide  :  expressing  this  proportion 
atomically,  the  proportion  of  carbon  to  oxygen 
is  12  to  32. 

The  gravimetric  composition  of  carbon 
dioxide  may  be  directly  determined  by  the 
combustion  of  a  known  weight  of  pure  carbon 
in  a  stream  of  oxygen  gas,  and  absorbing  and 
weighing  the  carbon  dioxide  that  is  formed. 
This  was  done  with  great  care  and  accuracy 
by  Dumas  and  Stas  in  the  experiments  by 
which  they  determined  the  atomic  weight  of 
carbon.  Fig.  68  represents  the  apparatus  em- 
ployed for  this  purpose.  A  weighed  quantity 


FIG.  67. 


of  diamond,  contained  in  a  small  platinum  boat,  was  introduced  into 
a  porcelain  tube,  which  could  be  strongly  heated  in  a  furnace.  The 
oxygen  for  its  combustion  was  contained  in  a  glass  bottle,  from 
which  it  could  be  expelled  by  allowing  water  to  enter  through  the 
funnel.  As  it  was  necessary  that  the  oxygen  should  be  absolutely 
free  from  any  carbon  dioxide,  the  water  used  in  the  little  gas- 
holder contained  potassium  hydroxide  in  solution.  The  oxygen 
was  then  passed  through  the  tubes  A,  B,  C,  in  order  to  deprive  it 
of  carbon  dioxide  and  moisture,  and  lastly  through  a  small  desic- 
cating tube,  d,  which  was  weighed  before  and  after  the  experiment. 
The  pure  dry  oxygen  then  entered  the  strongly  heated  tube,  and 
the  carbon  there  burnt  away  to  carbon  dioxide,  leaving  a  minute 
quantity  of  ash,  which  was  carefully  weighed  at  the  conclusion  of 


Carbon  "Dioxide 

the  experiment.  A  small  layer  of  copper  oxide  was  placed 
tube,  in  the  position  indicated 
in  the  figure,  in  order  to  oxidise 
any  traces  of  carbon  monoxide 
which  were  liable  to  be  formed 
into  the  dioxide.  The  product 
of  the  combustion  was  carried 
forward  by  the  stream  of  oxygen 
through  a  series  of  tubes  ;  d' 
is  a  small  weighed  desiccating 
tube,  the  weight  of  which,  if 
the  diamond  used  contained  no 
hydrogen,  should  remain  un- 
changed. It  then  passes  through 
the  bulbs  F  and  G,  where  the 
carbon  dioxide  is  entirely  ab- 
sorbed. To  arrest  aqueous 
vapour,  which  would  be  carried 
away  from  the  solution  in  these 
bulbs  by  the  escaping  oxygen, 
the  gas  is  passed  through  H, 
containing  fragments  of  solid 
potassium  hydroxide  ;  this  tube 
is  weighed  along  with  the  potash 
bulbs.  K  is  a  guard  tube  con- 
taining fragments  of  solid  potas- 
sium hydroxide,  in  order  to 
prevent  atmospheric  carbon 
dioxide  and  moisture  from  gain- 
ing access  to  the  weighed  por- 
tions of  the  apparatus. 

The  weight  of  the  diamond 
minus  the  weight  of  the  ash 
which  was  left  gave  the  actual 
weight  of  the  carbon  burnt  ; 
the  increase  in  weight  of  the 
tubes  gave  the  weight  of  the 
carbon  dioxide  which  was  pro- 
duced, and  this  weight,  minus 
the  weight  of  carbon  used,  gave 
the  weight  of  oxygen  that  was 


309 
in  the 


3IO  Inorganic  Chemistry 

consumed.  As  a  mean  of  a  number  of  experiments,  Dumas  and 
Stas  found  that  80  parts  of  oxygen  by  weight  combined  with  29.99 
parts  of  carbon. 

From  a  knowledge  of  the  density  of  carbon  dioxide  and  the 
volume  of  oxygen  it  contains,  we  know  that  the  molecule  of  this 
gas  contains  two  atoms  ;  therefore,  by  the  simple  equation — 

80  :  32  :  :  29.99  :  11.99, 

11.99  parts  of  carbon  combine  with  32  parts  of  oxygen,  and  the 
number  1 1.99  is  therefore  the  atomic  weight  of  carbon  as  determined 
by  these  chemists. 

The  Carbonates.— Although  carbonic  acid,  H2CO3,  is  a  very 
unstable  compound,  the  salts  it  forms  are  stable.  Being  a  dibasic 
acid,  it  is  capable  of  forming  salts  in  which  either  one  or  both  of 
the  hydrogen  atoms  have  been  replaced  by  an  equivalent  of  a 
metal  ;  thus  in  the  case  of  sodium  we  have — 

(1)  Disodium  carbonate  (normal  sodium  carbonate)          .  Na2CO3. 

(2)  Hydrogen  sodium  carbonate  (bicarbonate  of  soda)      .   HNaCO3. 

Similarly,  with  the  divalent  metal  calcium,  it  is  possible  to  form — • 

(1)  Normal  calcium  carbonate  ......  CaCO3, 

and — 

(2)  Hydrogen  calcium  carbonate  (bicarbonate  of  lime)     .  CaH2(CO3)2. 

The  formation  of  carbonates  by  the  action  of  carbon  dioxide  upon 
the  hydroxides  may  be  illustrated  by  the  following  equations  : — 

2KHO  +  CO2=K2CO3+  H2O. 
CaH2O2  +  CO2  =  CaCO3  +  H2O. 

The  first  of  these  changes  is  the  one  that  takes  place  when 
carbon  dioxide  is  absorbed  by  the  potassium  hydroxide  employed 
by  Dumas  and  Stas  in  the  course  of  their  experiments  already 
described.  The  second  equation  represents  the  reaction  which 
results  when  carbon  dioxide  is  passed  into  lime  water.  In  this 
latter  case,  if  the  gas  be  passed  through  the  turbid  solution  for 
some  time,  the  turbidity  will  gradually  disappear,  and  the  solution 
once  more  become  clear.  The  normal  calcium  carbonate  (CaCO3) 
which  is  first  formed,  and  which  is  insoluble,  is  converted  into 
the  soluble  bicarbonate,  CaH2(CO3)2.  If  this  solution  be  boiled, 
this  unstable  salt  is  decomposed  with  the  evolution  of  carbon 


Carbonates  311 

dioxide  and  water  and  the  reprecipitation  of  the  normal  calcium 
carbonate  — 

CaH2(CO3)2  =  CaCO3  +  H2O  +  CO2. 

The  presence  of  this  compound  in  natural  waters  is  associated 
with  the  property  known  as  the  hardness  of  water  (see  Natural 
Waters,  p.  221). 

When  one  volume  of  dry  carbon  dioxide  is  mixed  with  two  volumes  of  dry 
ammonia,  the  two  gases  unite,  forming  a  compound  known  as  ammonium 
carbamate  — 


XT  TT  \ 

CO2  +  2NH3=CO2,2NH3  or  ^^Q  j-CO, 


which  is  the  ammonium  salt  of  the  unknown  carbamic  acid,   j^Q2  i  CO. 

The  relation  between  this  compound  and  carbamide  or  urea  will  be  obvious 
by  an  inspection  of  the  formula  j^H  |  CO. 

This  substance  was  the  first  "  organic  "  compound  which  was  ever  obtained 
from  purely  inorganic  sources  (page  295).  It  can  be  obtained  by  the  action  of 
carbonyl  chloride  upon  ammonia— 

COC12+4NH3=CO(NH2)2+2NH4C1. 

Carbon  Suboxide,  C3O2.  —  By  the  action  of  phosphoric  acid  upon  ethyl- 
malonate  under  reduced  pressure,  and  at  a  temperature  about  300°,  a  mixture 
of  ethylene  and  carbon  suboxide  is  obtained  *  — 


The  products  of  the  action  are  condensed  in  a  vessel  cooled  by  liquid  air. 
The  ethylene  is  then  allowed  to  vaporise  at  the  ordinary  temperature,  leaving 
the  liquid  suboxide  behind.  This  is  purified  by  volatilisation  and  recon- 
densing  the  vapour  in  a  tube  cooled  to  about  -65°. 

The  compound  is  a  colourless  mobile  liquid  having  a  powerful  acrid  smell. 
It  boils  at  7°,  and  the  vapour  burns  with  a  blue  but  smoky  flame,  yielding 
carbon  dioxide. 

The  substance  may  be  regarded  as  the  anhydride  of  malonic  acid.  It 
readily  combines  with  water,  forming  malonic  acid  — 


o    = 
O  :  C  :  C  :  C  :  O  +  2H2O=HOOC  •  CH2  •  COOH. 


Diels  and  Wolf.     Ber.  ,  1906. 


CHAPTER  X 
COMPOUNDS  OF  CARBON  WITH  HYDROGEN 

THESE  two  elements  unite  together  in  various  proportions,  form- 
ing an  enormous  number  of  compounds  known  generally  under 
the  name  of  the  hydrocarbons.  The  reason  for  the  existence  of  so 
great  a  number  of  compounds  of  these  two  elements  is  to  be  found 
in  the  fact  that  the  atoms  of  carbon  possess,  in  a  very  high  degree, 
the  property  of  uniting  amongst  themselves.  This  property  of 
carbon  gives  rise  to  the  formation  of  a  number  of  groups  or  series 
of  compounds  the  members  of  which  are  related  to  each  other 
and  to  the  simplest  member  of  the  series.  Thus  the  compound 
methane,  CH4,  is  the  simplest  member,  or  the  "  foundation-stone," 
of  a  series  of  hydrocarbons  of  which  the  following  are  the  first 
four  : — 


Methane    .     .     .     CH4 
Ethane ....     C2H6 


Propane  .     .     .     C3H8 
Butane     .     .     .     C4H10 


It  will  at  once  be  seen  that  each  compound  differs  in  composi- 
tion from  its  predecessor  by  an  increment  of  CH2,  and  that  each 
may  be  expressed  by  the  general  formula,  CnH2n+2. 

Such  a  series  of  compounds  is  known  as  a  homologous  series,  and 
any  one  member  is  called  a  homologue  of  any  other. 

In  the  following  chapter  the  three  hydrocarbons,  methane, 
ethylene,  and  acetylene,  will  be  briefly  studied.  Each  of  these  is 
a  "  foundation-stone,"  or  starting-point,  of  a  series  similar  to  the 
one  already  mentioned  ;  thus — 

Methane,    CH4,    first  member  of  the  CnH2n+2  series  of  hydrocarbons. 
Ethylene,    C2H4,         „  ,,  CnH211 

Acetylene,  C2H2         ,,  ,,          CnH2n-2 

METHANE  (Marsh  Gas—Fire-Damp}. 
Formula,  CH4.     Molecular  weight =16.4.     Density =8.2. 

Occurrence.— Methane  is  found  in  the  free  state  in  large  quan- 
tities in  nature.  It  is  one  of  the  products  of  the  decompositions 

312 


Methane 


313 


which  has  resulted  in  the  formation  of  the  coal-measures.  It  is 
therefore  found  in  enormous  quantities  in  coal  mines,  where  it 
not  only  occurs  in  vast  pent-up  volumes,  under  great  pressure, 
which  escape  with  a  rushing  sound  when  the  coal  is  being  hewn  ; 
but  it  is  also  occluded  within  the  pores  of  the  coal.  Methane  is 
also  evolved  from  petroleum  springs. 

The  name  marsh  gas  has  been  given  to  this  compound,  on 
account  of  its  occurrence  in  marshy  places  by  the  decomposition  of 
vegetable  matter.  The  bubbles  of  gas  which  rise  to  the  surface 
when  the  mud  at  the  bottom  of  a  pond  is  gently  disturbed  consist 
largely  of  marsh  gas. 

Modes  of  Formation. — (i.)  When  a  mixture  of  sodium  acetate 
and  sodium  hydroxide  is  strongly  heated 
in   a   copper  retort,  sodium  carbonate  is 
produced  and  marsh  gas  is  evolved — 

CH3'COONa  +  NaHO  =  Na2CO3  +  CH4. 

The  gas  obtained  by  this  reaction  always 
contains  more  or  less  hydrogen. 

(2.)  Pure  methane  may  be  obtained  by 
the  decomposition  of  zinc  methyl,  by  means 
of  water — 


(3.)  The  most  convenient  method  for 
preparing  methane  is  by  the  action  of  zinc- 
copper  couple  upon  methyl  iodide.*  For 

this  purpose  the  zinc-copper  couple  is  placed  in  a  small  flask,  and 
a  mixture  of  equal  volumes  of  methyl  iodide  and  methyl  alcohol  is 
introduced  by  means  of  the  stoppered  funnel  (Fig.  69).  The  gas 
is  caused  to  pass  through  a  tube  filled  with  the  zinc-copper  couple, 
whereby  it  is  deprived  of  any  vapour  of  the  volatile  methyl  iodide, 
and  is  collected  over  water  in  the  pneumatic  trough. 

The  reaction  which  takes  place  is  essentially  a  reduction  of  the 
iodide  by  means  of  the  nascent  hydrogen  produced  by  the  action 
of  the  zinc-copper  couple  upon  the  alcohol  or  the  water  present, 
and  may  therefore  be  represented  by  the  equation — 

CH3I+2H  =  CH4+HI. 


*  "  Chemical  Lecture  Experiments,"  new  ed. ,  No.  449. 


314  Inorganic  Chemistry 

The  hydriodic  acid  must  not  be  regarded  as  escaping  as  such, 
but  in  the  presence  of  the  zinc  forming  a  compound  with  it.  If  water 
only  is  present,  the  compound  Znl'HO  is  formed  ;  while  if  methyl 
alcohol  is  employed  the  zinc  compound  will  have  the  composition 
ZnI*CH3O  :  the  complete  equation  (omitting  the  copper  which  does 
not  enter  into  the  chemical  change)  being  — 


Marsh  gas  is  formed  during  the  process  of  the  distillation  of  coal,  and  is 
therefore  a  large  constituent  of  coal  gas,  the  amount  varying  from  35  to  40 
per  cent. 

Properties.  —  Methane  is  a  colourless  gas,  having  no  taste  or 
smell.  It  burns  with  a  pale,  feebly  luminous  flame.  When  mixed 
with  air  or  oxygen  and  ignited  the  mixture  explodes  with  violence. 
The  products  of  its  combustion  are  water  and  carbon  dioxide, 
the  methane  requiring  twice  its  own  volume  of  oxygen  for  its 
complete  combustion,  and  yielding  its  own  volume  of  carbon 
dioxide  — 


Methane  is  only  about  one-half  as  heavy  as  air,  its  specific 
gravity  being  0.55  (air=  i).  The  fire-damp  of  coal  mines  is  nearly 
pure  methane,  its  average  composition  being  — 

Methane       .        .         .         .         .         .     96.0 

Carbon  dioxide    .         .         .         .         .0.5 

Nitrogen      ......       3.5 

100.0 

ETHYLENE  (Olefiant  Gas). 
Formula,  C2H4.     Molecular  weight=28.4.     Density=i4.2. 

Modes  of  Formation.  —  (i.)  This  compound  is  obtained  when 
ethyl  iodide  is  acted  upon  by  an  alcoholic  solution  of  potassium 
hydroxide  — 


(2.)  It  is  also  formed  when  ethylene  dibromide  is  brought  in  con- 
tact with  zinc-copper  couple,  the  ethylene  dibromide  being  diluted 
with  its  own  volume  of  alcohol  — 


(3.)  Ethylene  may  be  prepared   by  acting   upon   alcohol   with 
certain  powerful    dehydrating  agents,   such  as   phosphoric   pent- 


Ethylene  315 

oxide  or  sulphuric  acid,  the  latter  being  most  commonly  em- 
ployed. The  mixture  of  alcohol  and  sulphuric  acid  is  heated  in 
a  flask  to  about  165°,  when  a  brisk  effervescence  takes  place.  From 
the  point  of  view  of  the  final  products,  the  reaction  may  be  re- 
garded as  the  abstraction  of  the  elements  of  water  from  the 
alcohol,*  thus — 

C2H6O  —  H2O  =  C2ri4. 

1'he  action,  however,  is  always  accompanied  by  secondary  re- 
actions, which  result  in  the  rapid  blackening  of  the  mixture  owing 
to  the  separation  of  carbon.  The  sulphuric  acid  then  acts  upon 
this  carbon  with  the  evolution  of  carbon  dioxide  and  sulphur 
dioxide.  Hence  the  ethylene  that  is  obtained  by  this  process  is 
always  contaminated  with  considerable  quantities  of  these  gases, 
from  which  it  must  be  purified  by  being  passed  through  a  solution 
of  sodium  hydroxide. 

(4.)  Pure  ethylene  is  most  readily  prepared  by  the  action  of 
syrupy  phosphoric  acid  (the  ordinary  tribasic  acid)  upon  alcohol.t 
About  50  or  60  c.c.  of  the  acid  are  placed  in  a  small  Wurtz  flask 
of  about  1 80  c.c.  capacity.  The  flask  is  fitted  with  a  cork  carrying 
a  thermometer  and  a  dropping-tube  (Fig.  70),  the  end  of  the  latter 
being  drawn  out  to  a  fine  point  and  reaching  to  the  bottom  of  the 
flask.  The  acid  is  boiled  for  a  few  minutes  until  the  concentration 
reaches  such  a  point  that  the  temperature  rises  to  200°,  when  the 
alcohol  is  allowed  to  enter  drop  by  drop;  the  rate  at  which  the  alcohol 
is  admitted  being  visible  in  the  dropping-bulb.  By  keeping  the  tem- 
perature between  200°  and  220°,  a  steady  and  continuous  stream 
of  gas  is  evolved;  which  after  being  deprived  of  the  small  quantities 
of  ether  and  undecomposed  alcohol  with  which  it  is  accompanied, 
by  being  passed  through  a  small  Woulf's  bottle  standing  in  ice, 
is  practically  pure  ethylene.  The  action  of  the  phosphoric  acid  is 
the  same  as  that  of  sulphuric  acid,  the  first  action  being  the 
formation  of  phosphovinic  acid,  which  is  subsequently  decomposed 
in  a  similar  manner  to  the  sulphovinic  acid. 

*  In  reality  the  action  is  more  complex,  and  takes  place  in  two  stages,  the 
first  being  the  formation  of  ethyl  hydrogen  sulphate  or  sulphovinic  acid, 
C2H5'HSO4— a  compound  which  is  analogous  to  hydrogen  potassium  sulphate, 
KHSO4;  and  the  second  being  the  decomposition  of  this  compound  when 
heated  either  alone  or  in  the  presence  of  sulphuric  acid — 
C2H5-OH  +  H2SO4= C2H3  -HSO4  +  H2O, 

C2Hg-HS04  =  CaH4+  H2S04. 
t  Newth,  Jour.  Chem.  Soc.,  1901. 


Inorganic  Chemistry 

Properties. — Ethylene  is  a  colourless  gas,  having  a  somewhat 
pleasant  ethereal  smell  ;  it  burns  with  a  highly  luminous  flame, 
forming  carbon  dioxide  and  water,  one  volume  of  the  gas  requiring 
three  volumes  of  oxygen  for  its  complete  combustion,  and  produc- 
ing twice  its  own  volume  of  carbon  dioxide — 


C2H4 


2  =  2CO2  +  2H2O. 


If  mixed  with  oxygen  in  this  proportion  and  inflamed,  the  mix- 
ture explodes  with  great  violence. 


FIG.  70. 

When  mixed  with  twice  its  volume  of  chlorine  and  ignited,  the 
mixture  burns  rapidly  with  a  lurid  flame,  with  the  formation  of 
hydrochloric  acid  and  deposition  of  carbon — 


Acetylene  317 

Ethylene  is  rapidly  absorbed  by  fuming  sulphuric  acid  (more 
slowly  by  the  ordinary  strong  acid),  forming  ethyl  hydrogen 
sulphate — 

C2H4  +  H2SO4=C2H5'HSO4, 

and  from  this  compound,  by  distillation  with  water,  alcohol  may 
be  produced — 

C2H5'HSO4  +  H2O  =  C2H5-OH  +  H2SO4. 

Ethylene  is  reduced  to  the  liquid  state,  at  a  temperature  of  o°, 
by  a  pressure  of  41  atmospheres  ;  the  critical  temperature  of 
the  gas  is  +  10.1°,  at  which  point  a  pressure  of  51  atmospheres  is 
required  to  liquefy  it.  Liquefied  ethylene  boils  at  — 103°,  and  by 
increasing  its  rate  of  evaporation  temperatures  as  low  as  —140° 
can  readily  be  obtained.  Ethylene  (together  with  higher  members 
of  the  same  series)  constitutes  the  chief  illuminating  constituent  of 
ordinary  coal  gas,  of  which  it  forms  from  4  to  10  per  cent. 

ACETYLENE. 

Formula,  C2H2.     Molecular  weight =26. 2.     Density=i3.i. 

Modes  of  Formation. — (i.)  Acetylene  is  capable  of  being  syn- 
thetically formed  by  the  direct  union  of  its  elements.  For  this 


FIG.  71. 

purpose  a  stream  of  hydrogen  is  passed  through  a  three-way  globe, 
in  which  an  electric  arc  is  burning  between  two  carbon  rods, 
arranged  as  seen  in  Fig.  71  (a  quantity  of  sand  being  placed  in 
the  globe  to  prevent  fracture  from  falling  fragments  of  red-hot 
carbon).  Under  these  circumstances  a  small  quantity  of  the 
carbon  and  hydrogen  unites  to  form  acetylene,  which  is  swept  out 
of  the  globe  by  the  current  of  hydrogen.* 

*  The  formation  of  acetylene  appears  to  be  a  secondary  result,  due  to  the 
high  temperature  decomposition  of  methane  which  is  first  produced  (Bone, 
Jour.  Chem.  Soc.t  1897). 


318 


Inorganic  Chemistry 


(2.)  Acetylene  may  be  obtained  by  the  action  of  alcoholic  potash 
upon  ethylene  dibromide.  Alcoholic  potash  is  heated  in  a  flask, 
and  ethylene  dibromide  dropped  upon  it  from  a  stoppered  funnel, 
when  the  following  reaction  takes  place  — 

CHBr  +  2KHO==2KBr  +  2H 


(3.)  Acetylene  is  formed  when  marsh  gas  or  coal  gas  is  burned 
with  an  insufficient  supply  of  air  for  complete  combustion  ;  thus, 
when  a  Bunsen  lamp  becomes  accidentally  ignited  at  the  base  of 
the  chimney,  the  peculiar  and  unpleasant  smell  that  is  perceived  is 
partly,  though  not  entirely,  due  to  the  formation  of  acetylene. 

The  formation  of  acetylene  by  the  imperfect  combustion  of  coal 
gas  is  readily  shown  by  causing  a  jet  of  air  to  burn  in  an  atmos- 


FIG.  72. 

phere  of  coal  gas,  and  aspirating  the  products  of  combustion 
through  a  cylinder  containing  an  ammoniacal  solution  of  cuprous 
chloride,  as  shown  in  Fig.  72.  The  acetylene  is  absorbed  by  the 
ammoniacal  cuprous  chloride,  forming  a  deep-red  coloured  com- 
pound known  as  cuprous  acetylide  — 


Cu2Cl2,2NH3  +  H2O 


2  =  2NH4C1  +  C2H2Cu2O.* 


When  this  compound  is  acted  upon  by  hydrochloric  acid,  it  is 
decomposed  with  the  evolution  of  acetylene,  thus  — 

C2H2Cu2O  +  2HC1  =  Cu2Cl2  +  H2O  +  C2H2. 

*  Keiser  has  shown  that,  when  perfectly  dry,  the  compound  loses  a  mole- 
cule of  water,  and  has  the  composition  C2Cu2,  and  not  C2H2Cu2O  (or 
H2O)  ;  in  fact,  that  the  compound  is  a  carbide  of  copper. 


Acetylene  319 

Formerly  this  method  was  commonly  practised  when  any  quantity 
of  acetylene  was  required. 

(4.)  For  all  practical  purposes  acetylene  is  now  always-  prepared 
by  the  action  of  water  upon  calcium  carbide.  The  carbide  may  be 
placed  in  a  dry  flask  furnished  with  a  dropping  funnel  and  de- 
livery tube,  and  on  gradually  admitting  water  drop  by  drop  a 
rapid  evolution  of  nearly  pure  acetylene  at  once  takes  place  — 

CaC2  +  2H2O  =  Ca(HO)2  +  C2H2. 

Properties.  —  Acetylene  is  a  colourless  gas  having  an  extremely 
offensive  smell,  which  rapidly  induces  headache  ;  when  inhaled  in 
an  undiluted  state  it  is  poisonous.  The  gas  burns  with  a  highly 
luminous  and  smoky  flame.  When  burnt  from  specially  constructed 
jets  it  gives  a  pure  white  light  of  great  intensity,  and  on  this 
account  is  a  most  important  illuminant.  Acetylene  is  present  in 
small  quantities  in  ordinary  coal  gas,  and  its  presence  may  be 
detected  by  the  formation  of  the  red  precipitate  of  cuprous  acety- 
lide  when  coal  gas  is  allowed  to  bubble  through  an  ammoniacal 
solution  of  cuprous  chloride.  This  reagent  furnishes  not  only  a 
ready  and  delicate  test  for  the  presence  of  acetylene,  but  also 
provides  a  means  of  removing  this  gas  from  admixture  with  other 
gases.  Thus,  in  the  synthetic  formation  described  above,  the 
gases  issuing  from  the  globe  are  passed  into  a  flask  containing 
this  solution,  which  immediately  absorbs  the  acetylene.  When 
acetylene  is  subjected  by  prolonged  heating  to  a  temperature  short 
of  a  red  heat,  it  undergoes  polymerisation,  and  is  converted  into 
liquid  hydrocarbons,  of  which  benzene,  C6H6,  is  one. 

Nascent  hydrogen  converts  acetylene  into  ethylene  — 


From  acetylene,  therefore  (a  compound  which  can  be  syn- 
thetically prepared  from  its  elements,  carbon  and  hydrogen),  a 
great  number  of  "  organic"  compounds  can  be  built  up,  for,  as  has 
been  already  explained  (page  317),  from  ethylene  it  is  easy  to 
obtain  alcohol,  which  opens  the  door  to  the  preparation  of  a  vast 
number  of  other  organic  compounds. 

Coal  Gas.  —  When  coal  is  distilled,  the  volatile  products  obtained 
are  :  (i)  coal  tar  ;  (2)  an  aqueous  liquid  containing  ammonia  and 
other  products,  and  known  as  ammoniacal  liquor  j  (3)  coal  gas. 


320 


Inorganic  Chemistry 


Coal  gas,  after  being  subjected  to  ordinary  purification,  is  a 
mixture  of  gases  which  maybe  divided  into  three  classes,  namely  : 
illuminantS)  diluents,  and  impurities.  The  most  important  of 
these  substances  are — 

[  Ethylene,  C2H4 ;    propylene,  C3H6  ;  ~\ 

.  I       butylene,  C4HS      ....     (CnH2n)        I  About  6.5 

*'  |  Acetylene,  C2H2;  aUylene,  C3H4      .     (CnH2n_2)  (percent. 

I  Benzene,  C6H6 (CnH2n-6)  J 

Diluents. — Hydrogen,  marsh  gas,  carbon  monoxide       .     About  90  per  cent. 
Impurities. — Nitrogen,     carbon     dioxide,    sulphuretted 

hydrogen About  3.5  per  cent. 

The  composition  of  the  gas  is  largely  determined  by  the  nature 
of  the  coal  employed,  as  may  be  seen  from  the  following  analyses 
of  gas  from  bituminous  and  from  cannel  coal  :— 


From  Bituminous  Coal. 

From  Cannel  Coal. 

T   r\nAn*^ 

Manchester 

London  (Frankland). 

ijOnaon 
(Frankland). 

(Bunsen  and 
Roscoe). 

^1.24 

qcr.QA 

AC    Cg 

Marsh  gas  . 
Carbon  monoxide 

.      32-87 
.      12.89 

0*"       T" 

35-28 
7.40 

OO     .7T" 

41.99 
10.07 

4O-  5° 

34.90 
6.64 

Illuminants 

•        3.87 

3.56 

10.81 

6.46 

Nitrogen     .         .         . 

2.24 

2.46 

Carbon  dioxide  . 

0.30 

0.28 

1.19 

3.67 

Sulphuretted  hydrogen 

.         ... 

0.29 

CHAPTER    XI 
COMBUSTION 

WHEN  chemical  action  is  accompanied  by  light  and  heat,  the 
phenomenon  is  called  combustion.  All  exhibitions  of  light  and 
heat  are  not  necessarily  instances  of  combustion  ;  thus,  when  an 
electric  current  is  passed  through  a  spiral  of  platinum  wire,  or 
through  a  carbon  thread  in  a  vacuous  bulb  (as  in  the  familiar 
•"glow"  lamps),  these  substances  become  hot,  and  emit  a  bright 
Tight.  Neither  the  platinum  nor  the  carbon,  however,  is  under- 
going any  chemical  change,  and  therefore  the  phenomenon  is  not 
one  of  combustion.  The  materials  are  simply  being  heated  to  a 
state  of  incandescence  by  external  causes,  and  as  soon  as  these 
cease  to  operate,  the  glowing  substances  return  to  their  original 
condition  unchanged. 

Combustion  may  be  defined  as  the  chemical  union  of  two  sub- 
stances, taking  place  with  sufficient  energy  to  develop  light  and 
heat.  When  the  amount  of  light  and  heat  are  feeble,  the  combus- 
tion is  described  as  slow  or  incipient;  while,  on  the  other  hand, 
when  they  are  considerable,  the  combustion  is  said  to  be  rapid  or 
active.  The  true  nature  of  combustion  was  not  understood  until 
after  the  discovery  of  oxygen  in  1775.  From  about  the  year  1650 
until  after  that  important  discovery,  the  phlogistic  theory  was 
universally  adopted.  According  to  this  view,  a  combustible  body 
was  one  which  contained,  as  one  of  its  constituents,  a  substance  or 
principle  to  which  the  name  phlogiston  was  applied.  Easily  com- 
bustible substances  were  considered  to  be  rich  in  phlogiston,  while 
those  that  were  less  inflammable  were  held  to  contain  but  little  of 
this  ingredient.  The  act  of  combustion  was  regarded  as  the 
escape  of  this  principle  from  the  burning  substance.  Thus,  when 
a  metal  was  burnt  in  the  air,  it  was  considered  to  be  giving  off  its 
phlogiston,  and  the  material  that  was  left  after  the  combustion 
(which  we  now  know  to  be  the  oxide  of  the  metal)  was  regarded  as 
the  other  constituent  of  the  metal,  and  was  called  the  calx.  The 

321  X 


322  Inorganic  Chemistry 

metal,  therefore,  was  supposed  to  be  a  compound  of  a  calx  with 
phlogiston.  By  heating  a  calx  along  with  some  substance  rich  in 
phlogiston,  the  former  again  combined  with  this  principle  and  the 
metal  was  once  more  produced.  Thus,  when  the  calx  of  lead  was 
heated  with  charcoal  (a  substance  pre-eminently  rich  in  phlo- 
giston), the  charcoal  supplied  the  calx  with  the  necessary  amount  of 
phlogiston  to  produce  the  compound  of  calx  of  lead  and  phlogiston, 
which  was  metallic  lead.  This  theory  of  combustion,  after  sustain- 
ing many  severe  shocks  (from  such  experiments  as  those  of  Boyle 
and  others,  who  showed  that  the  calx  of  a  metal  was  heavier  than 
the  metal  used  in  its  formation),  received  its  death-blow  on  the 
discovery  of  the  compound  nature  of  water,  and  that  this  substance 
was  produced  by  the  combustion  of  hydrogen  in  oxygen. 

In  all  processes  of  combustion  it  is  customary  to  regard  one  of  the 
substances  taking  part  in  the  chemical  change  as  the  combustible, 
and  the  other  as  the  supporter  of  combustion.  Usually  that  sub- 
stance which  surrounds  or  envelops  the  other  is  called  the  sup- 
porter of  combustion.  Thus,  when  a  jet  of  burning  hydrogen  is 
introduced  into  a  jar  of  chlorine,  or  when  a  fragment  of  charcoal 
burns  in  oxygen,  the  chlorine  and  the  oxygen  are  spoken  of  as  the 
supporters  of  combustion,  while  the  hydrogen  and  carbon  are  termed 
the  combustibles. 

In  all  the  more  familiar  processes  of  combustion  the  atmosphere 
itself  is  the  enveloping  medium,  and  the  air  is  therefore,  par  excel- 
lence, the  supporter  of  combustion  ;  and  in  ordinary  language  the 
terms  combustible  and  incombustible  are  applied  to  denote  sub- 
stances which  burn,  or  do  not  burn,  in  the  air.  By  a  similar 
process  of  limitation,  it  has  become  customary  to  speak  of  other 
gases  as  supporters  or  non-supporters  of  combustion,  if  they  behave 
towards  ordinary  combustibles  as  air  does.  Thus  we  say  of  hydro- 
gen, or  marsh  gas,  or  coal  gas,  that  they  are  combustible,  but  do 
not  support  combustion ;  and  of  oxygen,  or  chlorine,  or  nitrous 
oxide,  that  they  do  not  burn,  but  will  support  combustion  ;  and 
lastly,  of  such  gases  as  ammonia,  or  carbon  dioxide,  or  sulphur 
dioxide,  that  they  neither  burn  nor  support  combustion. 

This  distinction,  however,  is  a  purely  conventional  one,  and  has 
little  or  no  scientific  significance ;  for,  by  a  slight  modification  of  the 
conditions,  either  hydrogen,  marsh  gas,  or  coal  gas  may  become 
supporters  of  combustion,  and  oxygen,  chlorine,  or  nitrous  oxide 
the  combustible  substances.  Thus,  when  a  jet  of  hydrogen  burns 
in  oxygen,  we  say  that  the  hydrogen  is  the  combustible,  and  the 


Combustion 


323 


oxygen  the  supporter  of  combustion  (Fig.  73,  A)  ;  but  if  a  jet  of 
oxygen  be  thrust  up  into  a  jar  of  hydrogen  (Fig.  73,  B),  it  ignites 
as  it  passes  the  burning  hydrogen,  and  continues  to  burn  in  the 
hydrogen. 

By  means  of  the  apparatus  shown  in  Fig.  74,  this  may  be  still 
more  strikingly  shown.*  A  stream  of  hydrogen  is  passed  into  the 
lamp  chimney  by  the  tube  H,  and  the  issuing  gas  inflamed  as  it 
escapes  at  the  top.  Oxygen  is  admitted  through  the  tube  O.  and 
the  jet  of  gas  ignited  by  pushing  the  long  tube  up  into  the  burning 


FIG.  73. 


FIG.  74. 


hydrogen  at  the  top,  and  then  drawing  it  down  to  the  position 
shown  in  the  figure,  where  the  jet  of  oxygen  continues  to  burn  in 
the  atmosphere  of  hydrogen. 

By  means  of  the  same  apparatus,  oxygen,  or  chlorine,  or  nitrous 
oxide  may  be  caused  to  burn  in  either  hydrogen,  marsh  gas,  or 
coal  gas.  Ammonia,  which,  as  already  mentioned,  is  usually 
described  as  being  neither  combustible  nor  a  supporter  of  com- 
bustion, when  surrounded  by  an  atmosphere  of  oxygen  is  readily 
inflammable,  and  will  as  readily  support  the  combustion  of  oxygen. 

The  atmosphere  itself  becomes  the  combustible  body  when  the 


"  Chemical  Lecture  Experiments,"  new  ed.,  No.  367. 


324  Inorganic  Chemistry 

usual  conditions  of  combustion  are  reversed.  Thus,  if  a  stream  of 
coal  gas  be  passed  through  a  similar  lamp  glass,  through  the  cork 
of  which  a  short  straight  glass  tube  passes  (Fig.  75),  air  will  be 
drawn  up  through  this  tube,  and  may  be  inflamed  by  passing  up  a 
lighted  taper.  The  jet  of  air  will  then  continue  to  burn  as  a  non- 
luminous  flame.  The  air  is  here  the  combustible,  and  the  coal  gas 
the  supporter  of  combustion.  If  the  excess  of  coal  gas  be  inflamed 
as  it  escapes  from  the  top,  the  opposite  conditions  will  be  fulfilled, 
the  air  being  the  supporter  of  combustion,  and  the  coal  gas  the 
combustible. 

This  interchangeableness  of  the  terms  combustible  and  sup- 
porter of  combustion  applies 
also  to  substances  that  are 
liquid  or  even  solid  at  the 
ordinary  temperature.  If  a 
small  quantity  of  some  inflam- 
mable liquid,  as  ether,  carbon 
disulphide,  turpentine,  &c.,  be 
boiled  in  a  flask,  and  the  issu- 
ing vapour  inflamed,  a  jet  of 
oxygen  gas  when  lowered  into 
the  flask  will  ignite  as  it  passes 
the  flame,  and  continue  to  burn 
in  the  vapour  of  the  liquid. 
In  the  same  way,  sulphur, 
which  is  a  combustible  solid, 
and  whose  vapour  is  inflam- 
mable in  the  air,  is  capable  in 
the  state  of  vapour  of  support- 
/  ing  the  combustion  of  oxygen. 


FIG.  75. 


— 5£'="  Since  combustion  is  the  result 
of  energetic  chemical  union, 
and  since  also  it  is  a  mere  condition  of  experiment  which  of  the 
two  acting  substances  shall  function  as  the  environment  of  the 
other,  it  will  be  seen  that  the  terms  "  combustible  "  and  "  supporter 
of  combustion,"  as  applied  to  a  chemical  substance,  do  not  express 
any  definite  or  characteristic  property  of  that  body. 

It  was  demonstrated  by  Boyle,  that  when  a  metal  is  burnt  in 
the  air,  the  calx  (or  oxide)  that  is  obtained  weighs  more  than  the 
metal  employed,  instead  of  less,  as  the  phlogistic  theory  seemed  to 
demand.  This  fact,  which  the  upholders  of  phlogiston  found  it  so 


Combustion 


.325 


difficult  to  reconcile,  is  seen  to  be  a  necessary  consequence  of 
combustion,  considered  from  the  modern  point  of  view.  In  all 
instances  of  combustion  the  weight  of  the  products  of  the  action 
is  equal  to  the  total  weight  of  each  of  the  two  substances  taking 
part  in  the  chemical  combination.  When,  for  example,  the  metal 
magnesium  burns  in  the  air,  the  weight  of  the  product  of  the  com- 
bustion is  equal  to  the  weight  of  the  metal,  plus  the  weight  of  a 
certain  amount  of  oxygen  with  which  it  united  in  the  act  of  burn- 
ing. This  gain  in  weight  during  combustion  may  be  demonstrated 
in  a  number  of  ways.  Thus,  if  a  small  heap  of  finely  divided  iron, 
obtained  by  the  reduction  of 
the  oxide,  be  counterpoised 
upon  the  pan  of  a  balance,  and 
then  ignited,  the  iron  will  be 
seen  to  burn,  and  as  it  burns 
the  balance  will  show  that  the 
smouldering  mass  is  increasing 
in  weight.  In  this  case  the 
sole  product  of  the  combustion 
is  a  solid  substance,  namely, 
iron  oxide,  which  remains  upon 
the  pan  of  the  balance  ;  but 
the  same  result  follows  when 
the  product  of  the  action  is 
gaseous.  Thus,  for  instance, 
when  a  fragment  of  sulphur  is 
burnt,  although  it  disappears 
from  sight,  it,  like  the  iron, 
combines  with  oxygen  to  form 
an  oxide.  This  oxide,  however,  FIG  ^ 

being  a  gas,  escapes  into  the 

atmosphere.  If  the  sulphur  be  burnt  in  such  a  manner  that  the 
sulphur  dioxide  is  collected  and  weighed,  it  also  will  be  found  to 
be  heavier  than  the  original  sulphur.  In  the  process  of  burning, 
i  gramme  of  sulphur  unites  with  about  I  gramme  of  oxygen,  and 
the  product  therefore  weighs  2  grammes.  By  causing  an  ordinary 
candle  to  burn  in  the  apparatus  shown  in  Fig.  76,  where  the  in- 
visible products  of  its  combustion  are  arrested,  the  increase  in 
weight  may  easily  be  seen.  The  candle  being  essentially  a  com- 
pound of  carbon  and  hydrogen,  the  products  of  its  burning  will  be 
carbon  dioxide  and  water,  both  of  which  will  be  absorbed  by  the 


326  Inorganic  Chemistry 

sodium  hydroxide  in  the  upper  part  of  the  tube.  Consequently, 
as  the  candle  burns  away,  the  arrangement  gradually  gains  in 
weight ;  the  increase  being  the  weight  of  the  atmospheric  oxygen 
which  has  combined  with  the  carbon  and  the  hydrogen  to  form 
the  compounds  carbon  dioxide  and  water. 

Heat  Of  Combustion. — During  the  process  of  combustion,  a 
certain  amount  of  heat  is  evolved,  and  a  certain  temperature  is 
attained — two  results  which  are  quite  distinct  The  temperatttre  is 
measured  by  thermometers  or  pyrometers,  while  the  amount  of 
heat  is  measured  in  terms  of  the  calorie,  or  heat  unit.* 

The  amount  of  heat  produced  by  the  combustion  of  any  sub- 
stance is  the  same,  whether  it  burns  rapidly  or  slowly,  provided 
always  that  the  same  final  products  are  formed  in  each  case. 
Thus,  when  I  gramme  of  phosphorus  burns  in  the  air  to  form 
phosphorus  pentoxide,  it  evolves  5747  calories  ;  and  when  the 
same  weight  of  phosphorus  is  burnt  in  oxygen,  although  the  com- 
bustion is  much  more  rapid  and  energetic,  and  the  temperature 
consequently  rises  higher,  the  amount  of  heat  evolved  is  precisely 
the  same. 

Again,  when  iron  is  heated  in  oxygen  it  burns  with  great  bril- 
liancy, and  with  evolution  of  much  heat ;  if,  however,  the  same 
weight  of  iron  be  allowed  slowly  to  combine  with  oxygen,  even 
without  any  manifestation  of  combustion,  it  is  found  that  the 
amount  of  heat  produced  in  forming  the  same  oxide  is  absolutely 
the  same. 

So  far,  therefore,  as  the  quantity  of  heat  produced  is  concerned, 
there  is  no  difference  between  active  combustion  and  slow  com- 
bustion, or  (confining  ourselves  to  the  case  of  combinations  with 
oxygen)  between  active  combustion  and  the  ordinary  process  of 
spontaneous  oxidation  at  ordinary  temperatures.  In  the  latter 
case  the  heat  is  given  out  slowly — so  slowly  that  it  is  conveyed 
away  by  conduction  and  radiation  as  fast  as  it  is  produced,  and 
consequently  the  temperature  of  the  material  undergoes  no  per- 
ceptible change.  In  the  case  of  active  combustion,  the  action  is 
crowded  into  a  few  minutes  or  seconds,  and,  as  all  the  heat  de- 
veloped is  evolved  in  this  short  space  of  time,  the  temperature 
of  the  substances  rapidly  rises  to  the  point  at  which  light  is 
emitted. 

That  heat  is  developed  during  the  process  of  spontaneous  oxida- 

*  The  major  calorie  sometimes  used  is  equal  to  1000  calories.  See  Thermo- 
chemistry, Part  I.  chap.  xv. 


Heat  of  Combustion  327 

tion  is  readily  shown.  Thus,  if  a  small  heap  of  fragments  of 
phosphorus  be  exposed  to  the  air,  it  will  be  evident  from  the 
formation  of  fumes  of  oxide  that  it  is  undergoing  oxidation.  As 
the  action  proceeds,  and  as  the  heat  produced  by  the  oxidation  is 
developed  more  rapidly  than  it  is  radiated  away  (especially  from 
the  interior  portions  of  the  heap),  it  will  be  seen  that  the  phos- 
phorus quickly  begins  to  melt,  and  finally  the  temperature  will 
rise  to  the  point  at  which  active  combustion  begins,  when  the  mass 
will  burst  into  flame. 

It  has  been  shown  that  many  destructive  fires  have  arisen  from 
masses  of  combustible  material,  such  as  heaps  of  oily  cotton  waste, 
undergoing  this  process  of  spontaneous  oxidation,  until  the  heat 
developed  within  the  mass  has  risen  sufficiently  high  to  inflame 
the  material.  To  the  operation  of  the  same  causes  is  to  be 
referred  the  spontaneous  firing  of  haystacks  which  have  been 
built  with  damp  hay,  and  also  the  spontaneous  inflammation  of 
coal  in  the  holds  of  ships. 

As  the  temperature  produced  by  combustion  is  augmented  by 
increasing  the  rapidity  with  which  the  chemical  action  takes  place, 
it  will  be  at  once  obvious  why  substances  which  burn  in  the  air, 
burn  with  increased  brilliancy  and  with  higher  temperature  in  pure 
oxygen.  In  the  air  every  molecule  of  oxygen  is  surrounded  by 
four  molecules  of  nitrogen,  therefore  for  every  one  molecule  of 
oxygen  that  comes  in  contact  with  the  burning  substance,  four 
molecules  of  this  inert  element  strike  it ;  and  by  so  doing  they  not 
only  prevent  the  contact  of  so  much  oxygen  in  a  given  interval  of 
time,  but  they  themselves  have  their  temperature  raised  at  the 
expense  of  the  heat  of  the  burning  material.  The  number  of 
oxygen  molecules  coming  in  contact  with  a  substance  burning  in 
the  air,  in  a  given  time,  may  be  increased  by  artificially  setting  the 
air  in  rapid  motion  :  hence  the  increased  rapidity  of  combustion 
(and  consequent  rise  of  temperature)  that  is  effected  by  the  use  of 
bellows,  or  by  increasing  the  draught  by  means  of  chimneys  and 
dampers. 

The  augmentation  of  temperature  obtained  by  the  substitution 
of  pure  oxygen  for  air  is  well  illustrated  in  the  case  of  burning 
hydrogen.  The  temperature  of  the  flame  of  hydrogen  burning 
in  oxygen,  known  as  the  oxy-hydrogen  flame,  is  extremely  high, 
and  when  allowed  to  impinge  upon  a  fragment  of  lime,  it  quickly 
raises  the  temperature  of  that  substance  to  an  intense  white  heat, 
when  it  emits  a  powerful  light— the  so-called  oxy-hydrogen  limelight. 


328 


Inorganic  Chemistry 


The  following  results  obtained  by  Bunsen  show  the  temperatures 
reached  by  the  combustion  of  hydrogen,  and  of  carbon  monoxide, 
in  air  and  in  oxygen — 

The  flame  of  hydrogen  burning  in  air     .         .         .  2024° 

oxygen     .         .  2844° 

,,  carbon  monoxide  burning  in  air        .  1997° 

»  »  »  oxygen  3003° 

It  will  be  seen  that  whereas  the  flame  of  hydrogen  in  air  is  hotter  than  that 
of  carbon  monoxide  in  air,  when  these  gases  burn  in  oxygen  the  temperature 


FIG.  77. 

of  the  carbon  monoxide  flame  is  higher  than  that  of  hydrogen.  This  is  due 
to  the  partial  dissociation  of  the  water  which  results  from  the  combustion  ot 
the  latter.  It  has  been  shown  that  when  a  mixture  of  hydrogen  and  oxygen,  in 
the  proportion  to  form  water,  is  ignited,  the  temperature  produced  by  the 
union  of  a  portion  of  the  mixture  rises  above  the  point  at  which  water  dis- 
sociates ;  and  consequently  for  a  certain  small  interval  of  time  a  condition  of 
equilibrium  obtains,  during  which  as  many  molecules  of  water  are  dissociated 
as  are  formed :  during  this  state  the  temperature  falls,  when  rapid  combus- 
tion once  more  proceeds.  It  will  be  seen,  therefore,  that  the  limits  to  the 
temperature  which  can  be  reached  by  combustion  are  influenced  by  the 
points  at  which  the  products  of  combustion  undergo  dissociation. 


Ignition  Point  329 

Ignition  Point. — The  temperature  to  which  a  substance  must 
be  raised  in  order  that  combustion  may  take  place  is  called  its 
ignition  point.  Every  combustible  substance  has  its  own  ignition 
temperature.  If  this  point  be  below  the  ordinary  temperature 
i.he  substance  will  obviously  take  fire  when  brought  into  the  air, 
without  the  application  of  heat ;  such  substances  are  said  to  be 
spontaneously  inflammable,  and  must  necessarily  be  preserved  out 
of  contact  with  air. 

Passing  from  cases  of  spontaneous  inflammability,  we  find  a 
very  wide  range  existing  between  the  igniting  points  of  different 
substances.  Thus,  a  jet  of  gaseous  phosphoretted  hydrogen  may 
be  ignited  by  causing  it  to  impinge  upon  a  test-tube  'containing 
boiling  water ;  carbon  disulphide  vapour  is  inflamed  by  a  glass 
rod  heated  to  120°,  while  the  diamond  requires  to  be  raised  nearly 
to  a  white  heat  before  combustion  begins. 

The  difference  between  the  temperatures  of  ignition  of  hydrogen 
and  marsh  gas  may  be 

well  seen  by  means  of  the  \ 

old  steel  mill  of  the  miner 
(Fig.  77).  By  causing  the  * 
steel  disk  to  revolve  at  a 
high  speed,  while  a  frag- 
ment of  flint  is  lightly 
pressed  against  its  edge,  a 
shower  of  sparks  isthrown 
out;  and  on  directing  a  jet  FIG.  78. 

of  hydrogen  upon  these 

sparks  the  gas  is  instantly  ignited,  while  they  may  be  projected 
into  a  stream  of  marsh  gas  without  causing  its  inflammation. 
The  same  fact  is  also  made  strikingly  apparent  by  depressing 
a  piece  of  fine  wire  gauze  upon  flames  of  marsh  gas  (or  coal 
gas)  and  hydrogen.  In  the  former  case  the  flame  will  not  pass 
through  the  gauze,  although  it  may  be  shown  that  marsh  gas 
is  making  its  way  through  by  applying  a  lighted  taper  imme- 
diately above  the  wire.  If  the  gauze  be  held  over  the  issuing  jet 
of  gas  the  latter  may  be  ignited  by  a  taper  upon  the  upper  side  of 
the  gauze,  but  the  combustion  will  not  be  communicated  to  the 
inflammable  gas  beneath  (Fig.  78).  The  gauze  conducts  the  heat 
away  from  the  flame  so  rapidly  that  ttie  temperature  of  the  metal 
does  not  rise  to  the  ignition  point  of  the  marsh  gas  on  the  other 
side,  and  therefore  the  combustion  cannot  be  propagated  through 


330  Inorganic  Chemistry 

the  gauze.  In  the  case  of  hydrogen,  however,  it  will  be  found  that 
the  instant  the  gas  upon  the  upper  side  of  the  gauze  is  inflamed 
the  flame  passes  through  and  ignites  the  hydrogen  beneath.* 

It  is  upon  this  principle  that  the  safety  of  the  "Davy  lamp"  depends. 
This  consists  of  an  ordinary  oil  lamp,  the  flame  of  which  is  surrounded  by 
a  cylinder  of  wire  gauze  (usually  made  double  at  the  top),  through  which  air 
to  supply  the  flame  freely  passes  in  and  the  products  of  combustion  pass  out. 
When  such  a  lamp  is  taken  into  an  atmosphere  in  which  marsh  gas  is  pre- 
sent, this  gas,  entering  through  the  gauze,  becomes  ignited  within  the  chimney, 
producing  a  very  characteristic  effect  upon  the  lamp  flame.  According  to  the 
amount  of  marsh  gas  present  the  flame  is  seen  to  become  more  and  more 
extended,  at  the  same  time  becoming  less  luminous,  until  the  whole  interior 
of  the  gauze  cylinder  is  filled  with  the  burning  gas,  emitting  a  faint  bluish 
light,  known  among  the  miners  as  \\\&  corpse- light.  The  burning  marsh  gas 
is  unable  to  communicate  its  combustion  to  the  inflammable  mixture  outside, 
for  the  same  reason  that  the  flame,  in  the  experiment  already  referred  to,  was 
unable  to  pass  through  the  wire  gauze.  If  from  any  cause  the  flame  should 
heat  any  spot  of  the  gauze  chimney  to  a  temperature  above  the  ignition  point 
of  marsh  gas,  the  outside  combustible  mixture  will  become  ignited.  It  has 
been  shown  that  by  exposing  the  lamp  to  a  strong  air  'draught  the  flame  may 
be  so  driven  against  the  gauze  as  to  unduly  heat  the  metal.  It  has  also  been 
proved  that  the  same  result  frequently  follows  from  the  explosive  wave  that 
is  produced  in  a  mine  when,  from  some  accidental  cause,  the  operation  of 
blasting  .(or  shot-firing]  results,  not  in  the  splitting  of  the  rock,  but  in  merely 
blowing  out  the  "  tamping."  The  violent  concussion  to  the  air  which  follows 
such  a  blown-out  shot  has  been  known  to  blow  the  flames  of  the  Davy  lamps, 
even  in  remote  parts  of  the  workings,  bodily  through  the  gauze ;  and  if  such 
lamps  are  burning  at  the  time  in  an  inflammable  mixture,  it  would  thereby  be 
fired. 

By  the  behaviour  of  the  flame  of  a  Davy  lamp  when  placed  into  an  atmos- 
phere containing  marsh  gas,  it  is  possible  to  estimate,  with  a  rough  degree  of 
accuracy,  the  percentage  amount  of  that  gas  which  is  present.  For  this  pur- 
pose the  flame  is  turned  down  as  low  as  possible,  and  the  height  to  which  the 
burning  marsh  gas  extends  (the  so-called  fire-damp  cap]  is  measured  against  a 
scale  graduated  in  tenths  of  inches.  Fig.  79  (two-thirds  the  actual  size)  shows 
the  "  caps  "  obtained  by  the  presence  of  4,  5,  and  6  per  cent,  of  marsh  gas.f 

When  the  ignition  point  of  a  substance  is  lower  than  the  tem- 
perature produced  by  its  combustion,  such  a  substance,  when 

*  Recent  experiments  of  Victor  Meyer  (Berichte,  No.  16,  1893),  upon  the 
ignition  temperature  of  explosive  gaseous  mixtures,  give  the  following  results : — 
A  mixture  of  oxygen  and  hydrogen  (electrolytic  gas)  explodes  at     612° 
Explosive  mixture  of  oxygen  and  marsh  gas      ....     656° 
,.  ,,  ,i  coal  gas         ....     647° 

f  In  a  recent  development  of  this  method  of  testing,  a  small  hydrogen 
flame  is  substituted  for  the  oil  lamp  flame,  whereby  it  is  possible  to  detect  the 
presence  of  0.25  per  cent,  of  marsh  gas  (Clowes). 


Ignition  Point  331 

ignited,  will  continue  to  burn  without  further  application  of  ex- 
ternal heat,  the  inflammation  being  propagated  from  particle  to 
particle  by  the  heat  developed  by  their  own  combustion.  All  the 
ordinary  processes  of  combustion  are  actions  of  this  order,  and 
belong  to  the  class  of  chemical  reactions  known  as  exothermic, 
that  is  to  say,  reactions  which  are  accompanied  by  an  evolution 
of  heat  (page  168). 

If,  on  the  other  hand,  the  ignition  point  be  higher  than  the  heat 
produced  by  chemical  union,  combustion  cannot  proceed  without 
the  continuous  application  of  external  heat.  The  igniting  point  of 


FIG.  79. 

nitrogen  in  oxygen,  for  example,  is  higher  than  the  temperature 
produced  by  the  union  of  these  elements  ;  therefore,  although  the 
nitrogen  may  be  ignited  by  the  heat  of  the  electric  spark,  it  is 
unable  to  communicate  its  combustion  to  contiguous  particles,  and 
the  inflammation  does  not  spread.  If  the  ignition  point  of  nitro- 
gen in  oxygen  had  been  lower  instead  of  higher  than  the  heat  of 
the  chemical  union  of  these  elements,  the  first  flash  of  lightning  that 
discharged  into  the  air  would  have  initiated  a  conflagration,  which 
would  have  extended  through  the  whole  atmosphere,  and  resulted 
in  the  removal  of  the  oxygen  and  its  replacement  by  oxides  of 
nitrogen. 


332  Inorganic  Chemistry 

The  production  of  acetylene  by  the  combination  of  carbon  with  hydrogen 
under  the  influence  of  high  temperature,  and  the  formation  of  cyanogen  and 
carbon  disulphide,  by  the  union  of  the  same  element  with  nitrogen  and  with 
sulphur  respectively,  are  illustrations  of  the  same  class  of  action  :  phenomena 
of  this  order  being  known  as  endothermic  reactions,  that  is,  reactions  that  are 
attended  with  an  absorption  of  heat  (page  168). 

Flame. — When  both  the  substances  taking  part  in  combustion 
are  gases  or  vapours,  the  sphere  of  the  chemical  action  assumes 
the  character  of  flame  ;  while,  on  the  other  hand,  if  one  of  the 
materials  is  a  solid  which  is  not  volatile  at  the  temperature  of  its 
combustion,  no  flame  accompanies  its  burning.  Such  solids  as 
sulphur,  phosphorus,  camphor,  wax,  &c.,  during  combustion  in  air, 
undergo  vaporisation,  and  consequently  burn  with  the  formation  of 
flame  ;  while  such  substances  as  iron,  copper,  carbon,*  &c.,  which 
do  not  pass  into  vapour  at  the  temperature  produced  by  their  com- 
bustion in  oxygen,  burn  in  this  gas  without  giving  rise  to  a  flame. 

Flames  differ  very  widely  in  their  general  appearance,  and  in 
the  majority  of  cases  are  distinctly  characteristic  :  thus,  hydrogen 
burns  in  air  with  a  flame  that  is  almost  absolutely  colourless,  and 
is  scarcely  visible  in  bright  daylight ;  sulphur  burning  in  air  pro- 
duces a  pale  blue  flame  ;  ammonia  in  oxygen  a  flame  having  a 
yellow-ochre  colour ;  carbon  monoxide  a  rich  blue  flame  ;  while 
cyanogen  burns  with  a  flame  having  the  delicate  'colour  of  the 
peach  blossom.  Other  flames  are  characterised  by  their  luminosity. 
Thus,  phosphorus  burning  in  oxygen  emits  a  dazzling  yellow  light, 
that  is  almost  blinding  to  the  eyes  ;  magnesium  burns  in  the  air 
with  an  intense  bluish-white  light ;  the  flame  produced  by  the 
combustion  of  the  vapour  of  nickel  carbonyl  in  air  emits  a  bright 
white  light  ;  and  the  flames  that  are  produced  by  most  hydro- 
carbons during  their  combustion  give  a  characteristic  yellowish- 
white  light. 

The  General  Structure  of  Flame. — The  simplest  form  of 
flame  is  one  that  is  obtained  by  the  combustion  of  a  substance 
which  itself  undergoes  no  decomposition,  and  in  which  the  product 
of  combustion  is  arrived  at  in  a  single  stage.  Such  flames,  for 
example,  as  that  of  hydrogen  burning  in  chlorine  or  in  air,  or  of 
carbon  monoxide  burning  in  air.  In  the  case  of  hydrogen  burning 
in  air,  the  materials  taking  part  in  the  process  being  elementary 

*  Under  certain  conditions  the  combustion  of  carbon  in  oxygen  is  accom- 
panied by  flame  ;  but  it  has  been  shown  that  at  the  temperature  at  which  this 
occurs  carbon  monoxide  is  being  formed. 


Flame 


333 


bodies,  no  complications  arising  from  decomposition  are  possible  ; 
and  although  carbon  monoxide  is  a  compound,  it  unites  with 
oxygen  without  itself  undergoing  any  decomposition,  and  passes 
directly  into  carbon  dioxide.  Such  flames  as  these,  when  burning 
from  the  end  of  a  tube,  consist  of  a  single  hollow  conical  sheath 
of  actively  burning  gas.  Fig.  80  represents  a  flame  of  burning 
hydrogen  :  the  darker  region  d  is  the  hollow  space  within  the  flame, 
consisting  of  unburnt  hydrogen  ;  while  the  flame  proper,  the  actual 
burning  portion,  is  the  sheath  b,  which  appears  practically  uniform 
throughout.  That  the  flame-cone  is  hollow  may  be  proved  by  a 
variety  of  experiments.  Thus,  if  a  sheet  of  white  paper  be  quickly 
depressed  into  a  flame,  a  charred  impression  of  the  section  of  the 
cone  will  be  obtained,  as  shown  in  Fig.  81,  from  which  it  will  be 


FIG.  80. 


FIG.  81. 


seen  that  no  combustion  is  taking  place  within  the  cone.  In  the 
same  way,  an  ordinary  lucifer  match  may  be  suspended  within  the 
flame,  where  it  will  remain  without  ignition  so  long  as  the  burning 
walls  of  the  flame  do  not  touch  it.  The  shape  of  a  flame  is  due  to 
the  fact,  that  as  the  gas  issues,  the  layer  nearest  to  the  walls  of  the 
tube  burn  round  the  orifice  of  the  tube  as  a  ring,  consequently  the 
next  layer  has  to  reach  up  above  this  ring  before  it  can  meet  with 
air  for  its  combustion,  and  each  successive  layer  has  to  pass  up 
higher  and  higher  in  order  to  find  its  supply  of  air,  and  in  this  way 
the  burning  area  is  built  up  into  the  form  of  a  cone.  To  show  that 
the  hollow  space  consists  of  unburnt  gas,  it  is  only  necessary  to 
insert  a  tube  into  the  interior  of  the  flame  in  such  a  way  as  to 


334 


Inorganic  Chemistry 


draw  off  a  portion  of  the  gas,  when  it  will  be  found  that  the  gas  so 
withdrawn  will  burn. 

Passing  from  this  simplest  type  to  substances  that  undergo 
decomposition  during  combustion,  or  which  yield  the  final  product 
of  oxidation  by  successive  stages,  it  is  found  that  the  flames  they 
give  rise  to  are  less  simple  in  structure. 

As  illustrations  of  various  degrees  of  complexity,  the  following 
examples  may  be  mentioned  : — 

(i.)  Ammonia  burning  in  oxygen.     This  flame  (Fig.  82)  is  very 
characteristic,  and  on  inspection  it  is  at  once  obvious  that  it  has  a 
less  simple  structure  than  the  hydrogen  flame.     In  this  case  the 
inner  hollow  portion  d  is  surrounded  by  a  double  flame-cone,  the 
inner  cone  a  having  a  yellow-ochre  colour,  and 
the  outer  portion  b  possessing  a  much  paler  colour, 
and  tending  to  green.     During  the  combustion 
of  ammonia,  the  compound  undergoes  decomposi- 
tion into  nitrogen  and  hydrogen.      This  decom- 
position, which  begins   in    the  hollow  region  d, 
takes  place  mainly  in  the  inner  cone  a,  and  the 
hydrogen    which    escapes    combustion    in    this 
region  passes    to    the  outside,  and  there  burns, 
forming  the  outer  cone.     Probably  there  is  also 
a  partial  combustion  of  the  nitrogen. 

(2.)  Carbon  disulphide  burning  in  air.  This 
flame,  like  the  ammonia  flame,  consists  of  a 
double  flame-cone,  consisting  of  an  inner  lilac- 
coloured  cone,  surrounded  by  an  outer  region 
having  a  deeper  blue  colour.  During  combus- 
tion carbon  disulphide,  like  ammonia,  is  decomposed,  but  in  this 
case  not  only  are  both  of  the  constituents  readily  combustible,  but 
the  carbon  passes  into  its  final  state  of  oxidation  in  two  stages, 
forming  first  carbon  monoxide  and  afterwards  carbon  dioxide. 

(3.)  Hydrocarbons  burning  in  air.  The  flames  produced  by  the 
combustion  of  these  compounds  include  those  which  are  commonly 
employed  for  illuminating  purposes,  such  as  candle,  gas,  and  oil 
flames,  and  in  all  essential  points  of  construction  they  are  practi- 
cally identical.  This  may  be  seen  to  be  the  case  by  a  comparison 
of  the  flames  of  a  candle  and  of  coal  gas  (Figs.  83  and  84).  In 
these  flames,  as  in  the  former  cases,  there  is  the  dark  hollow  space 
d,  consisting  of  heated  unburnt  gas  (in  the  candle  flame  this  gas 
is  generated  by  the  vaporisation  of  the  materials  of  the  candle, 


FIG.  82. 


Flame 


335 


which  in  the  melted  condition  are  drawn  up  the  wick  by  capillary 
action).  Above  this  there  is  a  region,  «,  which,  in  comparison 
with  the  rest  of  the  flame,  appears  almost  opaque,  and  which 
emits  a  bright  yellow  light.  This  luminous  area  constitutes  rela- 
tively the  largest  part  of  the  flame,  and  in  flames  that  are  used  for 
light-giving  purposes  it  is  intentionally  made  as  large  as  possible 
by  means  of  various  devices.  At  the  base  of  the  flame  there  is 
a  small  region,  c,  which  appears  bright  blue  in  colour,  and  is  non- 


FIG.  83. 


FIG,  84. 


luminous ;  and  surrounding  the  entire  flame  there  will  be  seen  a 
faintly  luminous  mantle,  b. 

The  flame  proper,  therefore,  consists  of  three  distinct  parts, 
namely  :  (i)  the  blue  region  c,  at  the  base  ;  (2)  the  faintly  luminous 
mantle  b  ;  and  (3)  the  yellow,  brightly  luminous  region  a.  These 
three  parts  constitute  the  flame-cone,  the  actual  area  of  combustion, 
which  envelops  the  dark  region  d ;  this,  as  already  stated,  consists 
of  unburnt  gas,  and  therefore  is  not,  strictly  speaking,  a  part  of  the 
flame. 

If  the  supply  of  gas  to  a  flame,  burning  as  represented  in  Fig.  84, 
be  diminished,  or  if  air  be  slowly  admitted  to  the  interior,  the  flame 


336  Inorganic  Chemistry 

will  shrink  down,  and  the  luminous  area  become  less  and  less, 
until  it  finally  disappears  altogether.  The  flame-cone  will  then  be 
found  to  consist  of  two  parts,  resembling  in  structure  the  double 
cone  of  the  ammonia  flame,  Fig.  82.  The  blue  region  c,  Fig.  84, 
which  is  only  fragmentary  in  the  flame  as  there  represented,  will 
have  become  continuous,  and  now  constitutes  the  inner  cone  ; 
while  the  mantle  b  forms  the  outer  cone,  the  flame  presenting  the 
appearance  seen  in  Fig.  85.  The  region  d^  as  before,  consists  of 
unburnt  gas. 

It  has  been  shown,  in  the  case  of  coal  gas  flames  burning  in  this 
manner,  that  in  the  inner  cone  c,  the  changes  going  on  result 
mainly  in  the  formation  of  carbon  monoxide  and  water,  together 
with  small  quantities  of  carbon  dioxide  and 
hydrogen  ;  and  that  in  the  outer  cone,  or 
mantle,  the  carbon  monoxide  and  hydrogen 
are  burning  to  carbon  dioxide  and  water. 
In  the  inner  cone,  therefore,  the  carbon  is 
burnt  to  its  first  stage  of  oxidation,  and  a 
portion  of  the  hydrogen  is  oxidised  to  water  ; 
in  the  outer  cone,  the  second  stage  of  oxida- 
tion  °f  the  carbon  takes  place  by  the  com- 
bustion of  the  carbon  monoxide  to  carbon 
dioxide,  and  the  hydrogen  which  escapes  combustion  in  the  inner 
cone  is  also  burnt. 

It  has  been  known  since  the  time  of  Dalton,  that  when  certain 
hydrocarbons  are  burnt  with  an  insufficient  amount  of  oxygen  for 
the  complete  oxidation  of  both  the  hydrogen  and  carbon,  carbon 
monoxide,  water,  and  hydrogen  are  produced.  This  result  is  pro- 
bably due  to  a  secondary  reaction  ;  the  first  stage  being  the  com- 
bustion of  hydrogen  to  form  water,  which  at  the  high  temperature 
is  then  decomposed,  either  by  the  carbon  or  the  hydrocarbons, 
according  to  the  following  equations  — 


The  various  parts  of  an  ordinary  gas  or  candle  flame,  therefore, 
are  due  to  the  different  chemical  reactions  that  are  taking  place  in 
these  areas  ;  these  changes  are  not  of  such  a  nature  that  they  can 
in  all  cases  be  perfectly  traced,  neither  is  one  set  of  reactions 
exclusively  confined  to  each  area,  but  rather  is  it  the  case  that 


Flame  337 

certain  cnemical  actions  predominate  in  each  particular  part  of  the 
flame. 

In  the  blue  region  c,  Figs.  83  and  84,  the  main  reactions  going 
forward  are  those  already  indicated,  by  which  carbon  monoxide, 
water,  and  hydrogen  are  produced.  In  the  faintly  luminous 
mantle  b,  carbon  monoxide  and  hydrogen  are  burning,  together 
with  small  quantities  of  hydrocarbons  which  may  have  escaped 
combustion  and  decomposition  in  the  luminous  region.  The  non- 
luminous  character  of  this  mantle  is  due  to  the  cooling  effect  of  the 
air  which  is  drawn  into  the  flame,  and  which  even  extinguishes 
combustion  upon  the  outer  limits  of  the  flame  before  every  trace  of 
combustible  material  is  burnt  ;  for  it  has  been  shown  that  small 
quantities  of  carbon  monoxide,  marsh  gas,  and  even  hydrogen 
escape  unburnt  from  a  gas  flame. 

The  chemical  decompositions  which  go  on  in  the  luminous  area 
cannot  be  said  to  have  been  thoroughly  established.  It  has  been 
shown  that  very  early  in  its  passage  up  the  flame  a  certain  amount 
of  the  marsh  gas  and  ethylene  present  is  converted  into  acetylene, 
the  change  taking  place  as  the  result  of  heat  alone.  The  gases 
ascending  the  dark  region  d  are  surrounded  on  all  sides  by  a  wall 
of  burning  material,  and  are  thereby  raised  in  temperature  to  the 
point  at  which  the  marsh  gas  and  ethylene  suffer  decomposition 
into  acetylene  and  hydrogen — 

2CH4=C2H2+3H2. 

The  following  table  (Lewes)  shows  the  gradual  development  of 
acetylene  in  such  a  flame  : — 

Total  Unsaturated 

Hydrocarbons.  Acetylene. 

Per  Cent.  Per  Cent. 

Gas  in  burner      ....     4.38  a°35 

\  inch  above  rim  of  burner.         .     4.00  0-340 

1 1  inch  above  rim        .         .         .1-53  0.560 

Tip  of  dark  region       .         .         .     1.98  1.410 

Centre  of  luminous  area      .         .     0.45  0.045 

Tip  of  luminous  area  .         .         .     o.oo  o.oo 

Therefore,  by  the  time  the  gases  have  reached  the  tip  of  the  dark 
region,  the  effect  of  heat  upon  them  has  been  to  raise  the  amount 
of  acetylene  to  over  70  per  cent,  of  the  total  unsaturated  hydro- 
carbons present.  As  the  acetylene  and  other  hydrocarbons  pass 
on  through  the  flame  along  with  steam,  carbon  dioxide,  and 

Y 


338  Inorganic  Chemistry 

carbon  monoxide,  other  and  more  complex  changes  go  on  whereby 
denser  hydrocarbons  are  formed,  and  carbon  itself  is  precipitated. 

The  formation  of  acetylene  in  that  region  of  the  flame  where  the 
coal  gas  is  in  excess  is  well  exemplified  in  the  case  of  air  burning 
in  an  atmosphere  of  coal  gas  (see  Fig.  75).  In  this  flame  the  air 
is  in  the  inside  and  the  coal  gas  upon  the  outside  ;  it  is,  in  effect,  an 
ordinary  coal  gas  flame  turned  inside  out.  The  formation  of  acety- 
lene, instead  of  taking  place  within  the  flame  (in  which  case  it  has 
to  pass  through  the  heated  area  and  is  thereby  decomposed), 
takes  place  upon  the  outer  surface  or  periphery  of  the  flame,  and 
therefore  largely  escapes  combustion  and  decomposition  and  passes 
away  into  the  coal  gas  atmosphere.  (See  Acetylene,  where  this 
method  is  described  for  the  preparation  of  this  compound.) 

The  Cause  of  Luminosity  in  Flames. — The  light-giving  property  of  a  flame 
is  not  due  to  the  operation  of  any  one  simple  cause.  It  was  at  one  time  sup- 
posed that  the  luminosity  of  a  flame  depended  solely  upon  the  presence  in  it 
of  suspended  solid  matter  resulting  from  the  chemical  decompositions  going 
on  during  combustion.  It  has  been  shown,  however,  that  this  general  state- 
ment does  not  satisfy  all  cases,  as  there  are  a  number  of  highly  luminous 
flames  in  which,  from  the  known  properties  of  the  products  of  combustion, 
there  cannot  possibly  be  any  solid  matter  present.  Thus,  for  example, 
phosphorus  burning  in  air  gives  a  flame  of  a  high  degree  of  luminosity  ;  but 
the  phosphorus  pentoxide  which  is  the  product  of  combustion,  although  solid 
at  ordinary  temperatures,  is  volatile  at  a  temperature  far  below  that  of  the 
flame.  The  same  may  be  said  of  the  luminous  flame  of  arsenic  burning  in 
oxygen,  where  the  still  more  volatile  arsenious  oxide  is  the  product. 

When  carbon  disulphide  burns  in  oxygen  or  in  nitric  oxide,  a  well-known  and 
intensely  luminous  flame  is  obtained,  in  which  only  gaseous  products  of  com- 
bustion can  be  present ;  and,  lastly,  the  flame  of  hydrogen  burning  in  oxygen 
can  be  made  under  certain  circumstances  to  emit  a  bright  light :  thus,  when  a 
mixture  of  these  gases  is  ignited  in  a  closed  eudiometer,  their  combustion  is 
attended  with  a  brilliant  flash  of  light,  the  only  product  being  water. 

There  are  three  causes  which  may  operate,  either  separately  or  together,  in 
imparting  luminosity  to  a  flame  or  in  increasing  its  light-giving  power  :  these 
are — (i.)  The  temperature  of  the  flame  ;  (2.)  the  density  of  the  flame  gases;  and 
(3.)  the  introduction  into  the  flame  of  solid  matter.  These  three  causes  will 
be  treated  separately  and  illustrations  given,  which,  so  far  as  our  knowledge 
extends,  can  be  directly  traced  to  the  independent  operation  of  each. 

(i.)  The  effect  of  temperature. 

(a.)  Upon  flames  in  which  solid  matter  is  known  to  be  absent. 

When  phosphorus  is  introduced  into  chlorine,  it  spontaneously  inflames  and 
burns  with  a  flame  of  such  extremely  feeble  luminosity  that  it  may  be  regarded 
as  non-luminous  ;  if,  however,  the  chlorine  be  previously  strongly  heated  by 
being  passed  through  a  red-hot  tube,  and  the  phosphorus  be  boilrng  when  it 
comes  in  contact  with  the  gas,  the  combustion  thus  started  upon  a  higher 


The  Luminosity  of  Flames  339 

platform    of  temperature   is   accompanied   by  a  flame  of  very  considerable 
luminosity. 

The  flame  of  carbon  disulphide  burning  in  air  emits  but  a  feeble  light ;  but 
when  this  substance  burns  in  pure  oxygen,  its  temperature  of  combustion  is 
greatly  raised  and  the  luminosity  of  the  flame  is  enormously  increased. 

Phosphoretted  hydrogen  burning  in  air  gives  a  flame  of  considerable  lumi- 
nosity ;  but  when  this  flame  is  fed  with  pure  oxygen,  and  its  temperature 
thereby  raised,  it  becomes  intensely  luminous. 

(/3.)  Upon  flames  in  which  solid  matter  is  known  to  be  present. 

The  flames  produced  by  the  combustion  of  zinc  or  magnesium  in  the  air, 
and  in  which  the  solid  oxides  are  present,  have  their  luminosity  greatly  in- . 
creased  when  pure  oxygen  is  substituted  for  air  and  the  temperature  of  com- 
bustion thereby  augmented. 

The  same  result  is  seen  in  the  case  of  flames  in  which  the  solid  matter  is 
artificially  introduced,  as  in  the  familiar  Welsbach  burner,  where  a  solid  gauze 
mantle,  composed  of  an  alkaline  earth,  is  placed  in  the  flame-cone  of  a  non- 
luminous  gas  flame,  thereby  rendering  it  luminous.  If  the  temperature  of 
this  flame  be  augmented  by  feeding  it  with  oxygen,  the  light  emitted  by  the 
incandescent  solid  is  greatly  increased. 

(7.)  Upon  flames  in  which  solid  matter  is  believed  to  be  present,  such  as 
candle,  gas,  and  other  hydrocarbon  flames. 

When  a  candle  or  gas  flame  is  introduced  into  oxygen,  although  it  shrinks 
in  size,  its  luminosity  is  increased.  It  has  also  been  shown  that  when  a  coal 
gas  flame  is  chilled  by  causing  it  to  spread  against  a  cold  surface,  its  luminosity 
is  diminished  or  destroyed  altogether ;  and,  conversely,  if  the  gas  and  the  air 
supplying  the  flame  be  strongly  heated  before  combustion,  the  luminosity  is 
greatly  increased.  In  this  case,  however,  the  direct  effect  of  change  of  tem- 
perature is  complicated  by  the  decompositions  going  on  in  the  flame  ;  for,  as 
already  mentioned,  the  conversion  of  the  non-illuminating  marsh  gas  into  the 
highly  illuminating  gas  acetylene  is  a  function  of  the  temperature. 

The  increase  of  light  obtained  from  a  gas  flame  by  previously  heating 
the  gas  and  air  is  the  principle  underlying  all  the  so-called  recuperative 
burners. 

It  is  evident,  therefore,  that  most  flames  gain  luminosity  by  having  their 
temperature  raised.  There  are,  however,  cases  in  which  increase  of  tempera- 
ture alone  appears  to  exert  no  influence  upon  the  luminosity.  The  flame  of 
hydrogen,  for  example,  which  is  practically  non-luminous  when  burning  in 
air,  does  not  become  more  luminous  when  burnt  in  oxygen,  although  its 
temperature  is  greatly  increased. 

(2. )  The  influence  of  the  density  of  the  flame  gases. 

It  has  been  shown  by  Frankland  *  that  the  luminosity  of  flame  is  intimately 
associated  with  the  pressure  to  which  it  is  subjected,  or  with  the  density  of  the 
flame  gases.  Thus,  it  is  found  that  a  gas  or  candle  flame,  when  burnt  either 
at  high  altitudes  cr  in  artificially  rarefied  atmospheres,  has  its  luminosity 
greatly  reduced  ;  and,  per  contra,  when  caused  to  burn  under  increased  pres- 
sure, the  luminosity  is  increased.  In  the  case  of  hydrocarbons,  complication 
arises  from  the  fact  that  the  temperature  of  the  flame  is  changed  by  alterations 

*  Phil.  Trans.,  vol.  cli.  p.  629;  Proc.  Royal  Society,  vol.  xvi.  p.  419. 


340 


Inorganic  Chemistry 


of  pressure.  Under  diminished  pressure  the  temperature  falls,  and  although 
there  is  less  loss  of  heat  by  radiation  in  rarefied  air  than  in  air  at  the 'ordinary 
pressure,  it  is  possible  that  the  general  lowering  of  the  temperature  of  the 
flame  may  modify  the  chemical  decompositions  in  the  direction  already  re- 
ferred to. 

Flames  other  than  those  of  hydrocarbons,  however,  and  in  which  no  solid 
matter  can  exist,  are  found  to  become  luminous  when  the  density  of  the  flame 
gas  is  increased  by  pressure.  Thus,  the  flame  of  carbon  monoxide  in  oxygen 
at  ordinary  pressures  emits  a  moderate  light ;  but  when  exposed  to  a  pressure 
of  two  atmospheres  the  luminosity  is  greatly  increased.  Even  the  non-luminous 
flame  of  hydrogen  burning  in  oxygen  becomes  luminous  under  a  pressure  of 
two  atmospheres,  and  when  examined  by  the  spectroscope  is  found  to  give  a 


FIG.  86. 


FIG.  87. 


continuous  spectrum.  It  has  been  found,  as  a  general  rule,  that  dense  gases 
and  vapours,  when  heated,  become  incandescent  or  luminous  at  much  lower 
temperatures  than  those  of  low  specific  gravity  ;  thus,  if  different  gases  be 
raised  to  incandescence  by  the  passage  through  them  of  electric  sparks,  under 
similar  conditions,  it  is  seen  that  the  light  emitted  by  the  glowing  vapour 
varies  with  the  density  of  the  gas.  The  luminosity  of  glowing  oxygen  (density, 
16)  is  greatly  superior  to  that  of  hydrogen  (density,  i),  while  the  light  emitted 
when  the  sparks  are  passed  through  chlorine  (density,  35.5)  is  considerably  in 
advance  of  either.  And  it  is  found  that  in  one  and  the  same  gas  the  luminosity 
of  the  spark  increases  as  the  density  is  increased  by  artificial  compression. 
Other  things  being  equal,  it  may  be  said  that  the  denser  the  vapours  which  are 
present  the  more  luminous  is  the  flame. 
(3.)  The  introduction  of  solid  matter  into  flames. 


The  Bunsen  Flame  341 

Non-luminous  flames  may  be  rendered  luminous  by  the  intentional  introduc- 
tion into  them  of  solid  matter,  which,  by  being  raised  to  a  sufficiently  high 
temperature,  will  become  strongly  incandescent.  Thus,  the  ordinary  lime- 
light owes  its  luminosity  to  the  incandescence  of  the  fragment  of  lime,  which 
is  raised  to  a  bright  white  heat  by  the  high  temperature  of  the  non-luminous 
oxy-hydrogen  flame.  The  lime  is  not  vaporised  at  the  temperature  of  the 
flame,  the  light  being  entirely  due  to  the  glowing  solid  matter. 

The  "  Welsbach  "  burner,  already  referred  to,  is  another  example  of  the  same 
order,  the  luminosity  in  this  case  being  due  to  the  introduction  into  an 
ordinary  non-luminous  Bunsen  flame  of  a  fine  gauze  mantle  made  of  thoria 
or  other  metallic  oxide  (Fig.  86).  When  such  a  mantle  is  raised  to  incandes- 
cence by  the  heat  of  the  gas  flame,  it  emits  a  bright  white  light,  strongly 
resembling  that  of  an  ordinary  Argand  gas  flame.  A  flame  may  also  be 
rendered  luminous  by  the  intentional  precipitation  within  it  of  carbon,  which, 
by  its  ignition  and  its  combustion,  produces  a  high  degree  of  luminosity.  Thus, 
if  a  small  quantity  of  alcohol  be  boiled  in  a  flask,  and  a  jet  from  which  chlorine 
is  issuing  be  then  lowered  through  the  burning  vapour  into  the  flask,  as  shown 
in  Fig.  87,  the  chlorine  will,  burn  in  the  alcohol  vapour  with  a  luminous  flame  ; 
and  the  precipitated  carbon  (which  is  thrown  out  of  combination  by  the  action 
of  the  chlorine  upon  the  alcohol),  ascending  into  the  previously  non-luminous 
alcohol  flame,  will  render  it  brightly  luminous. 

From  these  considerations  it  will  be  evident  that  the  luminosity  of  a  flame 
may  be  due,  first,  to  the  presence  of  vapours  sufficiently  dense  to  become 
incandescent  at  the  temperature  of  the  flame ;  or,  second,  to  the  presence  of 
solids  rendered  incandescent,  either  by  the  heat  of  the  flame  gases  alone,  or 
in  conjunction  with  their  own  combustion  ;  or,  third,  from  the  simultaneous 
operation  of  all  these  causes.  Ordinary  gas  and  candle  flames  come  under 
the  last  of  these  heads.  The  decompositions  that  go  forward  in  these  flames 
not  only  give  rise  to  dense  vapours  which  become  incandescent,  but  also  to 
the  precipitation  of  solid  carbon,  which  by  its  ignition  and  combustion  adds 
to  the  luminosity  of  the  flame. 

The  Bunsen  Flame. — The  construction  of  the  Bunsen  lamp  is  too  well 
known  to  need  description.  The  gas,  issuing  from  a  small  jet  situated  at  the 
base  of  a  metal  tube,  and  mixing  with  air  which  is  drawn  in  through  openings  in 
the  tube,  burns  at  the  top  of  the  chimney  with  the  familiar  non-luminous  flame. 
The  existence  of  this  flame  in  its  ordinary  condition  depends  upon  two  main 
causes ;  first,  upon  the  fact  that  in  the  immediate  neighbourhood  of  a  jet  of 
gas  issuing  from  a  small  orifice,  there  is  a  reduction  of  pressure  ;  and,  second, 
upon  the  relation  between  the  velocity  at  which  the  gases  pass  up  the  tube 
and  the  rate  of  propagation  of  combustion  in  the  mixture  of  air  and  coal  gas. 
Upon  the  first  of  these  causes  depends  the  entrance  of  air  into  the  "  air-holes  " 
of  the  lamp,  and  upon  the  second  depends  the  continuance  of  the  flame  in  its 
position  upon  the  top  of  the  tube. 

As  the  coal  gas  issues  from  the  small  jet  at  the  base  of  the  chimney,  instead 
of  the  gas  escaping  through  the  side-holes,  air  is  drawn  into  the  tube  by  virtue 
of  the  reduced  pressure  produced  immediately  round  the  jet.  That  this  area  of 
reduced  pressure  actually  exists  in  the  neighbourhood  of  the  jet  of  a  Bunsen 
may  be  proved  by  attaching  a  delicate  manometer  to  the  air-hole  of  such  a 
lamp,  as  shown  in  Fig.  88.  As  the  gas  is  turned  on,  the  liquid  in  the  horizontal 


342  Inorganic  Chemistry 

tube  will  be  sucked  towards  the  lamp,  showing  that  the  issuing  gas  causes  a 
partial  vacuum  in  its  immediate  neighbourhood.* 

In  order  that  the  flame  shall  remain  at  the  top  of  the  tube,  there  must  be  a 
certain  relation  between  the  velocity  of  the  issuing  gases  and  the  rate  of  pro- 
pagation of  combustion  in  the  mixture ;  for  if  the  latter  be  greater  than  the 
former,  the  flame  will  travel  down  the  tube  and  ignite  the  gas  at  the  jet  below. 
By  gradually  reducing  the  supply  of  gas  to  the  flame,  and  so  altering  the  pro- 
portion of  gas  and  air  ascending  the  tube,  the  mixture  becomes  more  and 
more  explosive,  until  a  point  is  reached  when  the  velocity  of  inflammation  is 
greater  than  the  rate  of  efflux  of  the  gases,  and  the  flame  travels  down  the  tube, 
and  the  familiar  effect  of  the  flame  "  striking  down"  is  obtained. 

The  same  result  may  be  brought  about,  and  the  effect  more  closely  observed, 
by  extending  the  chimney  of  the  lamp  by  means  of  a  wide  glass  tube.     As  the 
supply  of  gas  is  reduced,  or  the  quantity  of  air  introduced  is  increased,  the 
flame  will  be  seen  to  shrink  in  size  and  finally  descend  the  tube.     By  adjust- 
ment it  may  be   caused  either  to  ex- 
plode rapidly  down   the   tube    or    to 
travel  quite  slowly,  or  even  to  remain 
stationary  at  some  point  in  the  tube, 
which  is  slightly  constricted,  and  where, 
therefore,  the  flow  of  the  issuing  gas  is 
slightly  accelerated,  f 

The  non-luminosity  of  a  Bunsen 
flame  is  due  to  the  combined  opera- 
tion of  three  causes,  namely,  oxidation, 
dilution,  and  cooling.  It  was  formerly 
supposed  that  the  destruction  of  the 
luminosity  of  a  gas  flame  by  the  ad- 
mixture of  air  with  the  gas  before 
burning  was  entirely  owing  to  the 
influence  of  the  oxygen  in  bringing 
FIG.  88.  about  a  more  rapid  and  complete 

state  of  oxidation,  that  the  hydro- 
carbons were  at  once  completely  burnt  up  by  the  additional  uupply  of  oxygen 
so  provided.  It  has  been  shown,  however,  that  not  only  is  this  effect  brought 
about  by  air,  but  also  by  the  use  of  such  inert  gases  as  nitrogen,  carbon  dioxide, 
and  even  steam.  The  following  table  (Lewes)  shows  the  relative  volumes  of 
various  gases  that  are  required  to  destroy  the  luminosity  of  a  gas  flame  : — 

I  volume  of  coal  gas  requires  0.5  volumes  of  oxygen. 

,,  ,,  ,,  1.26          ,,        carbon  dioxide. 

,.  »  I.  2.27          ,,        air. 

,,  ,,  2.30          ,,         nitrogen. 

,,  ,,  ,,  5.11          ,,        carbon  monoxide. 

That  the  atmospheric  oxygen  effects  the  result  by  a  direct  oxidising  action, 
and  is  not  acting  merely  as  nitrogen  does,  is  proved  by  the  fact  that  mixtures 
of  oxygen  and  nitrogen,  containing  a  higher  proportion  of  oxygen  than  is 

*  See  "  Chemical  Lecture  Experiments,"  new  ed.,  498-502.       f  Ibid.,  506. 


The  Bunsen  Flame 


343 


present  in  air,  destroy  the  luminosity  more  rapidly  than  is  effected  by  air. 
Thus,  when  mixtures  containing  nitrogen  and  oxygen  in  the  proportion  of  3  to 
i,  2  to  i,  i  to  i  by  volume  are  employed,  the  volumes  of  the  mixtures  required 
to  destroy  the  luminosity  of  one  volume  of  coal  gas  are  respectively  2.02,  1.49, 
and  i. oo. 

It  has  been  shown  that  when  coal  gas  is  diluted  with  nitrogen  a  higher 
temperature  is  necessary  to  effect  its  decomposition ;  hence  the  action  of  the 
atmospheric  nitrogen  in  causing  the  loss  of  luminosity  of  a  gas  flame  is  in  part 
due  to  the  higher  temperature  that  is  required  for  the  formation  of  acetylene, 
which,  as  already  mentioned,  is  the  first  step  in  the  decomposition  and  con- 
densation of  the  hydrocarbons  in  the  gas. 

As  already  mentioned,  the  luminosity  of  a  flame  is  very  much  influenced  by 
alterations  of  temperature  ;  and  just  as  the  non-luminosity  of  the  outer  mantle 
of  an  ordinary  flame  is  partly  due  to  the  cooling  action  of  the  air  which  is 
dragged  into  the  flame  from  the  outside,  so  the  want  of  luminosity  of  the 
Bunsen  flame  is  in  part  due  to  the  cooling  influence  of  the  large  volume  of  air 
that  is  drawn  up  into  the  interior  of  the  flame.  That  the  gases  which  are 
drawn  into  a  flame  reduce  the  luminosity  by  virtue  of  their  cooling  action  is 
borne  out  by  the  fact  that  the  higher  the  specific  heat  of  the  diluent  (and 
therefore  the  greater  its  power  to  abstract  heat  from  the  flame)  the  less  of 
it  is  required  to  effect  the  destruction  of  the  luminosity  ;  thus,  as  already  men- 
tioned, less  carbon  dioxide  than  nitrogen  is  necessary  to  render  a  flame  non- 
luminous  :  the  specific  heat  of  nitrogen  is  0.2370,  while  that  of  carbon  dioxide 
is  0.3307. 

The  specific  heat  of  oxygen  is  also  slightly  greater  than  that  of  nitrogen, 
being  0.2405;  but  the  cooling  effect  of  dilution  with  this  gas  is  enormously 
overpowered  by  the  increased  temperature  due  to  its  oxidising  action  upon 
the  combustible  materials  of  the  flame. 

Experiments  made  upon  the  actual  temperatures  of  various  regions  of  a 
Bunsen  flame,  rendered  non-luminous  by  admixture  with  different  gases,  the 
results  of  which  are  seen  in  the  following  table  (Lewes),  show  the  cooling  effect 
of  these  diluents  upon  the  flame : — 

Temperature  of  Flame  from  Bunsen  Burner,  burning  6  cubic  feet  of  Coal 
Gas  per  Hour. 


Flame  rendered  Non- 

luminous  by 

Region  in  Flame. 

Air. 

Nitrogen. 

Carbon 
Dioxide. 

\  inch  above  burner   .                          . 

Degrees. 
135 

Degrees. 

54 

Degrees. 
3° 

Degrees. 
35 

i£  inch  above  burner  . 

421 

175 

III 

70 

Tip  of  inner  cone 

913 

1090 

444 

393 

Centre  of  outer  cone  . 

1328 

1533 

999 

770 

Tip  of  outer  cone 

728 

"75 

"51 

95  * 

Side  of  outer  cone,  level  with  tip  of  ) 
inner  cone                                           ) 

1236 

1333 

1236 

970 

L 

344  Inorganic  Chemistry 

In  the  case  of  air,  it  will  be  seen  that  the  first  effect  is  to  cool  the  flame  ; 
but  in  the  upper  region,  where  the  oxidising  action  of  the  oxygen  is  felt,  the 
temperature  rapidly  rises  to  a  maximum  at  a  point  about  half-way  between 
the  tip  of  the  inner  and  outer  cones.  In  the  flames  rendered  non-luminous 
by  the  two  inert  gases,  the  highest  temperature  is  only  reached  at  the  outer 
limit,  where  the  full  amount  of  oxygen  for  combustion  is  obtained  from  the 
outer  atmosphere. 

On  account  of  the  wide  range  of  temperature  exhibited  by  the  various 
regions  of  a  Bunsen  flame,  it  constitutes  a  most  valuable  analytical  instru- 
ment, for,  by  the  judicious  use  of  the  different  parts  of  the  flame,  it  is  often 
possible  to  detect  the  presence  of  several  flame-colouring  substances  in  a 
mixture.  Thus,  if  a  mixture  of  sodium  and  potassium  salts  be  introduced 
upon  platinum  wire  into  the  cooler  region  of  the  flame  near  its  base,  the  more 
volatile  potassium  compound  will  impart  its  characteristic  violet  tint  to  the 
flame  before  the  sodium  salt  is  volatilised  sufficiently  to  mask  the  colour,  by 
the  strong  yellow  it  itself  gives  to  the  flame.  In  this  way  many  mixtures  may 
readily  be  differentiated. 

If  a  piece  of  copper  wire  be  held  horizontally  across  a  Bunsen  flame,  so  as 
to  cut  the  inner  cone,  it  will  be  seen  that  the  wire  in  contact  with  the  edges 
of  the  flame  becomes  coated  with  copper  oxide,  while  the  portion  in  the  centre 
remains  bright.  On  moving  the  wire  so  as  to  bring  the  oxidised  portion  into 
the  inner  region,  the  oxide  will  be  reduced,  the  metal  once  more  becoming 
bright.  The  outer  area  of  a  flame,  where  oxygen  is  in  excess,  is  called  the 
oxidising  flame ;  while  the  inner  region,  in  which  heated  and  unburnt  hydro- 
gen or  hydrocarbons  exist,  is  spoken  of  as  the  reducing  flame.  These  regions 
exist  in  all  ordinary  flames.  The  oxidising  action  of  the  outer  flame  of  a 
candle,  for  example,  is  illustrated  in  the  behaviour  of  the  wick.  So  long  as 
the  wick  remains  in  the  inner  region  of  the  flame  it  is  not  burnt ;  and  in  the 
early  days  of  candles,  as  the  tallow  gradually  consumed,  the  wick  remained 
standing  straight  up,  and  by  degrees  extended  into  the  luminous  area  of  the 
flame,  where,  owing  to  the  deposition  of  soot  upon  it,  it  frequently  developed 
a  cauliflower-like  accretion,  which  greatly  impaired  the  luminosity  of  the 
flame,  and  which  necessitated  the  use  of  snuffers.  In  the  modern  candle, 
owing  to  a  method  of  plaiting  the  wick,  it  is  caused  to  bend  over  (as  shown 
in  Fig.  83),  and  so  thrusts  its  point  into  the  oxidising  region,  where  it  is 
continually  burnt  away. 


PART   III 

THE  SYSTEMATIC  STUDY  OP  THE  ELEMENTS, 
BASED  UPON  THE  PERIODIC  CLASSIFICA- 
TION. 


CHAPTER   I 


THE  ELEMENTS  OF  GROUP  VII.  (FAMILY  B.) 


Fluorine,  F   . 
Chlorine,  Cl  . 


19.00 
35-45 


Bromine,  Br 
Iodine,  I 


79-92 
126.92 


THE  first  to  be  discovered,  and  the  most  important  element  of 
the  group,  is  chlorine,  which  is  a  constituent  of  sea  salt  (sodium 
chloride).  The  term  halogen,  signifying  sea  salt  producer,  has 
been  applied  to  this  family  of  elements,  on  account  of  the  close 
resemblance  between  their  sodium  salts  and  sea  salt.  This  family 
exhibits,  in  a  marked  manner,  many  of  the  features  which  are 
found  to  exist  in  most  chemical  families  of  elements. 

In  their  general  behaviour  they  strongly  resemble  one  another, 
and  readily  displace  each  other  in  combinations  without  producing 
any  very  marked  change  upon  the  character  of  the  compounds. 
They  each  unite  with  hydrogen,  giving  rise  respectively  to  hydro- 
fluoric acid,  HF  ;  hydrochloric  acid,  HC1 ;  hydrobromic  acid, 
HBr;  hydriodic  acid,  HI* 

These  hydrogen  compounds  are  all  colourless  gases,  which  fume 
strongly  in  the  air ;  they  are  extremely  soluble  in  water,  and  are 
strongly  acid  in  character.  In  combination  with  potassium  and 
with  sodium,  the  halogens  form  a  series  of  compounds,  which  are 
similarly  constituted,  and  which  closely  resemble  each  other  in  their 

*  Some  chemists  name  these  compounds  hydrogen  fluoride,  hydrogen 
chloride,  hydrogen  bromide,  hydrogen  iodide  respectively,  and  employ  the 
names  hydrofluoric  acid,  hydrochloric  acid,  &c.,  to  denote  the  aqueous 
solutions  only. 

345 


346  Inorganic  Chemistry 

habits.     Their  similarity  of  composition  is  expressed  in  the  fol- 
lowing formulas  : — 

Compounds  with  potassium,  KF,  KC1,  KBr,  KI. 

Compounds  with  sodium,  NaF,  NaCl,  NaBr,  Nal. 

The  physical  properties  of  the  elements  exhibit  a  regular  grada- 
tion with  increasing  atomic  weight ;  thus,  fluorine  and  chlorine  are 
gases,  bromine  is  liquid,  while  iodine  is  solid  at  ordinary  tempera- 
tures. In  their  chemical  activity  they  also  show  the  same  gradual 
change;  thus,  in  the  case  of  their 'combination  with  hydrogen, 
when  fluorine  and  hydrogen  are  brought  together,  combination 
instantly  takes  place  with  explosion,  even  in  the  dark.  Chlorine 
and  hydrogen  do  not  combine  in  the  dark,  but  in  diffused  daylight 
they  unite  slowly,  and  in  direct  sunlight  their  combination  takes 
place  suddenly  with  explosion. 

Bromine  vapour  and  hydrogen  do  not  combine  even  in  direct 
sunlight,  but  a  mixture  of  the  two  gases  ignites  in  contact  with  a 
flame,  yielding  hydrobromic  acid,  while  iodine  vapour  and  hydro- 
gen require  to  be  strongly  heated  in  contact  with  spongy  platinum 
to  effect  their  combination.    This  difference  in  the  activity  of  the 
halogens  towards  hydrogen  is  seen  by  a  comparison  of  the  heats 
of  formation  of  their  hydrogen  compounds,  thus— 
H  +  F  =  HF  +38,500  cal. 
HC1  + 22,000   „ 
=  HBr+  8,440   „    * 
H  +  I    =HI    -   6,040    „    t 

Although  a  strong  resemblance  exists  between  all  the  members 
of  the  halogen  family,  the  element  fluorine,  which  is  the  typical 
member  (see  page  115),  stands  marked  off  from  the  others  in  many 
of  its  attributes.  Thus  fluorine  exhibits  a  great  tendency  to  form 
double  salts  which  have  no  counterpart  among  the  compounds  of 
the  other  elements  of  the  family,  and  at  temperatures  below  32° 
the  molecule  of  hydrofluoric  acid  consists  of  two  atoms  of  hydrogen 
and  two  of  fluorine,  having  the  composition  H2F2. 

FLUORINE. 

Symbol,  F.     Atomic  weight =19. 

History. — This  element,  the  first  of  the  halogen  series,  was  the 
most   recent   to   be  isolated,   it   having  baffled   all    attempts   to 
*  This  value  refers  to  bromine  in  the  liquid  state, 
f  Iodine  as  solid. 


Fluorine  347 

obtain  it  until  the  year  1886,  when  Moissan  succeeded  in  solving 
the  problem. 

Occurrence.  —  Fluorine  occurs  in  considerable  quantities  in  com- 
bination with  calcium  in  the  mineral  fluor  spar  (CaF2),  which  is 
found  in  cubical  crystals.  On  account  of  the  occurrence  of  this 
mineral  in  large  quantities  in  Derbyshire  it  is  frequently  termed 
Derbyshire  spar.  It  is  a  constituent  also  of  cryolite,  Na3AlF6,  fluor- 
apatite,  3P2O8Ca3,CaF2,  and  many  others.  In  small  quantities 
fluorine  is  found  in  bones,  in  the  enamel  of  teeth,  and  also  in 
certain  mineral  waters. 

Mode  of  Formation.—  When  an  electric  current  is  passed  into 
an  aqueous  solution  of  hydrochloric  acid,  the  acid  is  decomposed 
into  its  elements,  chlorine  being  liberated  at  the  positive  electrode, 
while  hydrogen  is  evolved  at  the  negative.  When  aqueous  hydro- 
fluoric acid  is  treated  in  the  same  way,  the  water  only  is  decom- 
posed, oxygen  and  hydrogen  being  liberated.  Davy  found  that  the 
more  nearly  the  acid  approached  the  anhydrous  condition,  the  less 
easily  did  it  conduct  electricity  ;  and  that  in  the  perfectly  pure 
state,  that  is,  entirely  free  from  water,  hydrofluoric  acid  was  a  non- 
conductor. Moissan's  recent  success  in  the  isolation  of  fluorine 
depends  upon  the  discovery  that  a  solution  of  the  acid  potassium 
fluoride,  HF,KF,  in  anhydrous  hydrofluoric  acid  is  an  electrolyte* 
and  that  by  the  passage  of  an  electric  current  through  this  solution 
fluorine  is  disengaged  at  the  anode,  or  positive  electrode,  and 
hydrogen  is  evolved  at  the  cathode. 

The  primary  products  of  the  electrolysis  are  potassium  (at  the 
cathode)  and  fluorine  at  the  anode.  The  potassium  then  reacts 
with  the  hydrofluoric  acid,  re-forming  potassium  fluoride  and 
liberating  an  equivalent  of  hydrogen  — 


Or,  expressed  in  the  form  of  ionic  equations  — 


The   reaction  is  performed  in  a   U-tube  made  of  an  alloy  of 
platinum  and  iridium,  a  material  which  is  less  acted  upon  by  the 


348 


Inorganic  Chemistry 


liberated  fluorine  than  platinum  alone.  The  apparatus  has  two 
side-tubes  (Fig.  89),  which  can  be  either  closed  with  a  screw  cap,  C, 
or  connected  to  platinum  delivery  tubes  by  means  of  the  union  D. 
The  two  limbs  of  the  tube  are  closed  by  means  of  stoppers  made 
of  fluor  spar,  shown  in  section  at  S,  and  which  can  be  securely 
screwed  into  the  tube.  These  serve  to  insulate  the  electrodes, 
which  are  constructed  of  the  same  platinum-iridium  alloy.  The 
anhydrous  hydrofluoric  acid  is  introduced  into  the  apparatus,  and 
about  25  per  cent,  of  its  weight  of  the  acid  potassium  fluoride  is 
added,  which  readily  dissolves  in  the  liquid.  The  tube  is  immersed 
in  a  bath  of  methyl  chloride  (M,  Fig.  90),  which  boils  at  -23°  ;  the 

supply  being  continuously  re- 
plenished from  the  reservoir  B, 
while  the  vapour  is  drawn  away 
by  the  pipe  C.  On  passing  a 
current  from  20  to  25  Grove's 
cells  through  the  apparatus, 
c  fluorine  is  evolved  at  the  posi- 
tive electrode,  and  hydrogen  is 
liberated  at  the  negative.* 

Properties — Fluorine  is,  of 
all  known  elements,  the  most 
chemically  active.  It  is  on 
account  of  its  intense  chemical 
affinities  that  it  so  long  resisted 
all  attempts  to  isolate  it,  as 
when  liberated  from  combina- 
tion it  instantly  combined  with 
FlG>  89-  the  materials  of  the  vessels  in 

which  the  reactions  were  made. 

It  is  impossible  to  collect  this  gas  by  any  of  the  usual  methods,  for 
it  decomposes  water  and  instantly  combines  with  mercury.  It  also 
attacks  glass,  so  that  it  can  only  be  collected  by  displacement  of  air 
in  vessels  of  platinum.  Fluorine  is  a  pale  yellowish-coloured  gas, 
appearing  almost  colourless  when  viewed  in  small  quantities.  The 
smell  of  the  gas  is  very  characteristic — it  is  irritating  to  the  mucous 
membranes,  and  is  not  unlike  the  odour  of  the  mixture  of  chlorine 
and  chlorine  peroxide,  evolved  from  potassium  chlorate  and  hydro- 

*  More  recently  Moissan  employs  a  copper  tube  of  300  c.c.  capacity,  fitted 
with  large  platinum  electrodes.  By  keeping  the  temperature  about  —50°, 
and  using  a  current  of  15  amperes,  he  obtains  the  gas  in  large  quantities. 


Fluorine 


349 


chloric  acid.  Whether  the  smell  actually  perceived  is  the  true 
smell  of  fluorine  is  doubtful,  for  when  fluorine  comes  into  contact 
with  the  moisture  in  the  nostrils  water  is  decomposed,  with  the  for- 
mation of  ozonised  oxygen  and  hydrofluoric  acid. 

Fluorine  not  only  decomposes  potassium  iodide,  with  liberation 
of  iodine,  but  also  displaces  chlorine  from  sodium  chloride. 

It  combines  directly  with  a  large  number  of  elements  with 
intense  energy  ;  in  contact  with  hydrogen  it  instantly  explodes. 
Iodine,  sulphur,  and  phosphorus  first  melt,  and  then  take  fire  in 
fluorine.  Crystals  of  silicon,  when  brought  into  the  gas,  spontane- 


llllllilililllllllllllllW  ' 

FIG.  90. 

ously  inflame,  and  burn  with  brilliancy.  All  of  the  metals  are  acted 
upon  by  fluorine  ;  some,  when  finely  divided,  undergoing  spontane- 
ous inflammation  when  thrown  into  the  gas.  Even  gold  and  plati- 
num are  attacked  by  fluorine,  especially  if  gently  warmed  ;  its  action 
upon  the  latter  metal  being  seen  by  the  corrosion  of  the  apparatus, 
and  especially  the  positive  electrode  employed  in  its  preparation. 

Organic  compounds  are  attacked  by  fluorine  with  violence,  and 
often  inflamed. 

When  fluorine  is  cooled  to  a  temperature  about  -  187°  (i.e.  a  few 


350  Inorganic  Chemistry 

degrees  below  the  temperature  of  boiling  oxygen,  obtained  by 
boiling  the  oxygen  under  slightly  reduced  pressure)  it  condenses 
to  the  liquid  state.*  Liquid  fluorine  is  a  mobile  yellow  liquid, 
resembling  liquid  chlorine.  Its  specific  gravity  is  1.14.  It  is 
without  action  upon  silicon,  phosphorus,  sulphur,  or  glass  ;  it  can 
therefore  be  produced  and  contained  in  glass  vessels.  Even  at 
this  low  temperature,  however,  fluorine  attacks  hydrogen  and 
hydrocarbons.  When  cooled  by  liquid  hydrogen  it  forms  a  pale 
yellow  solid,  melting  at  —223°.  On  cooling  the  solid  to  —252°  it 
loses  its  yellow  colour  and  appears  perfectly  white. 

HYDROFLUORIC  ACID  (Hydrogen  Fluoride}. 
Formula,  HF.     Molecular  weight=2o.oi.     Density  =10. 

Modes  Of  Formation.—  (  i.)  Hydrofluoric  acid  is  produced  when 
powdered  calcium  fluoride  (fluor  spar)  is  acted  upon  by  strong 
sulphuric  acid  — 


This  method  is  employed  for  the  commercial  preparation  of 
aqueous  solutions  of  hydrofluoric  acid.  The  mixture  of  fluor  spar 
and  sulphuric  acid  is  gently  warmed  in  a  leaden  retort,  and  the 
gaseous  acid  passed  into  water  contained  in  leaden  bottles.  This 
aqueous  acid  is  sent  into  the  market  in  gutta-percha  bottles. 

(2.)  The  anhydrous  acid  is  prepared  by  heating  hydrogen  potas- 
sium fluoride  (acid  potassium  fluoride)  in  a  platinum  retort.  The 
double  fluoride  of  potassium  and  hydrogen  splits  up  into  normal 
potassium  fluoride  and  hydrofluoric  acid  — 

HF,KF  =  KF  +  HF. 

For  this  purpose  the  perfectly  dry  double  fluoride  is  placed  in 
a  platinum  retort,  which  is  screwed  to  a  platinum  condensing 
arrangement,  as  seen  in  Fig.  91.  The  wooden  trough  through 
which  the  long  tube  passes  is  filled  with  a  freezing-mixture,  and 
the  platinum  bottle  is  also  surrounded  by  a  similar  mixture. 

Properties.  —  Anhydrous  hydrofluoric  acid  is  a  colourless, 
limpid,  strongly  fuming  liquid,  which  boils  at  19.5°.  It  has  a 
powerful  affinity  for  water,  and  can  only  be  preserved  in  perfectly 
stoppered  platinum  vessels,  which  are  kept  in  a  cool  place.  The 
acid  at  once  attacks  gutta-percha.  Gore  found  that  the  anhydrous 
acid  was  without  action  upon  glass. 

Pure  hydrofluoric  acid  is  an  extremely  dangerous  substance  to 
manipulate  ;  its  vapour,  even  when  diluted  with  air,  has  a  most 
*  Moissan,  May  1897. 


Hydrofluoric  Acid  351 

irritating  and  injurious  effect  upon  the  respiratory  organs,  and  if 
inhaled  in  the  pure  state  causes  death. 

A  single  drop  of  the  liquid  upon  the  skin  causes  the  most  painful 
ulcerated  sores,  accompanied  by  distressing  aching  pains  through- 
out the  whole  body.  The  metals  potassium  and  sodium  dissolve  in 
pure  hydrofluoric  acid,  with  the  formation  of  fluorides  and  evolution 
of  hydrogen. 

At  temperatures  above  88°  the  vapour-density  of  hydrofluoric 
acid  corresponds  to  the  formula  HF.  As  the  temperature  is 
lowered  the  molecules  aggregate  together,  and  the  density  of  the 
vapour  steadily  rises,  until  at  a  few  degrees  above  the  boiling- 


. 


FIG.  91. 


point  it  approaches  what  would  be  required  for  molecules  of  H3F3. 
At  about  32°  the  density  is  20  ;  but  whether  this  signifies  the  exis- 
tence of  molecules  having  the  composition  H2F2,  or  whether  it 
merely  represents  a  certain  mixture  of  more  complex  molecules, 
HnFn,  with  molecules  of  HF,  has  not  been  definitely  determined. 

Gaseous  hydrofluoric  acid  rapidly  attacks  glass,  and  it  is  largely 
employed  for  etching  purposes,  both  for  obtaining  designs  upon 
glass  and  for  the  purpose  of  etching  graduations  upon  glass  mea- 
suring instruments.  The  object  to  be  etched  is  first  coated  with 
wax,  and  the  design  or  other  marks  cut  upon  the  wax  by  means  of 
a  pointed  steel  tool.  In  this  way  the  surface  of  the  glass  is  laid 


352  Inorganic  Chemistry 

bare  in  parts,  and  on  exposing  the  object  to  the  action  of  the  acid, 
either  as  gas  or  aqueous  solution,  the  glass  is  rapidly  eaten  into, 
where  the  surface  has  been  exposed.  Its  action  upon  glass  is 
due  to  the  readiness  with  which  it  attacks  silicates,  the  fluorine 
combining  with  the  silicon  to  form  silicon  tetrafluoride  — 


Crystallised  silicon,  when  gently  heated,  takes  fire  in  gaseous 
hydrofluoric  acid,  giving  silicon  fluoride  and  hydrogen. 

Hydrofluoric  acid  is  extremely  soluble  in  water,  forming  a 
strongly  acid  corrosive  liquid,  which  readily  dissolves  many  of 
the  metals  with  evolution  of  hydrogen  — 


=  FeF2+H2. 
Silver  and  copper  are  also  dissolved  by  this  acid. 

CHLOEINE. 

Symbol,  Cl.     Atomic  weight =35. 45.     Molecular  weight =70. 90. 

History. — Chlorine  was  discovered  by  Scheele  (1774),  but  wa~ 
regarded  by  him  as  a  compound  substance.  He  applied  to  it  the 
name  of  dephlogisticaled  marine  acid  air,  having  obtained  it  by 
the  action  of  hydrochloric  acid  upon  ores  of  manganese.  The 
belief  that  chlorine  was  a  compound  of  oxygen  and  hydrochloric 
acid  was  generally  held  until  Davy's  time,  and  gave  rise  to  the 
name  of  oxymuriatic  acid. 

The  elementary  nature  of  chlorine  was  proved  by  Davy  (1810), 
who  gave  to  it  the  name  chlorine,  in  allusion  to  the  greenish-yellow 
colour  of  the  gas. 

Occurrence. — In  the  uncombined  condition  chlorine  does  not 
occur  in  nature.  In  combination  with  metals,  as  chlorides,  chlorine 
is  very  abundant,  the  commonest  chloride  being  sodium  chloride 
(common  salt). 

Many  of  the  salts  found  in  the  Stassfurt  deposits  consist  largely 
of  chlorides  (see  Alkali  Metals).  Chlorides  of  the  alkali  metals 
are  also  found  in  animal  secretions  and  in  certain  plants.  Chlorine 
occurs  in  combination  with  hydrogen,  as  hydrochloric  acid,  in 
volcanic  gases,  and  also  in  the  gastric  juice. 

Modes  Of  Formation.— ( i.)  When  hydrochloric  acid  is  poured 
upon  manganese  dioxide,  and  the  mixture  kept  cool,  a  dark-brown 


Chlorine 


353 


solution  is  obtained,  which  rapidly  decomposes  at  a  slight  rise  of 
temperature  with  the  evolution  of  chlorine. 

It  has  not  yet  been  clearly  established  whether  this  brown  solu- 
tion consists  of  the  compound  MnCl4  or  MnCl3,  formed  according 
to  one  of  the  equations — 


or — 


When  this  dark-brown  solution  is  gently  warmed,  the  higher 
chloride  breaks  up  into  manganous  chloride  (MnCl2)  and  chlorine  ; 
the  complete  reaction  being  expressed  by  the  equation  — 


The  experiment  is  conveniently  carried  out  in  the  apparatus  seen 
in  Fig.  92.     The  mixture  of  manganese  dioxide  and  hydrochloric 


FIG.  92. 

acid  is  gently  heated  in  a  large  flask,  and  the  gas,  after  being 
passed  through  water  in  the  Woulf's  bottle,  may  be  collected  by 
downward  displacement,  as  shown  in  the  figure.* 

(2.)  Instead  of  employing  hydrochloric  acid,  the  materials  from 
which  this  compound  is  prepared,  namely,  sodium  chloride  and 
stilphuric  acid,  may  be  used.  Thus,  if  a  mixture  of  sodium 


See  Experiment  154,  "  Chemical  Lecture  Experiments,"  new  ed. 

Z 


354  Inorganic  Chemistry 

chloride,  manganese  dioxide,  and  sulphuric  acid  be  gently  warmed, 
chlorine  is  readily  evolved  — 


It  will  be  seen  that  by  this  reaction  the  whole  of  the  chlorine  con- 
tained in  the  reacting  compounds  is  evolved  as  gas,  while  in  the 
former  case  a  part  of  it  .  remains  in  combination  with  the  man- 
ganese. 

(5.')  Many  other  highly  oxygenised  compounds,  when  acted  upon 
oy  hydrochloric  acid,  evolve  chlorine  ;  thus,  when  crystals  of  potas- 
sium dichromate  are  drenched  with  hydrochloric  acid  and  the 
mixture  heated,  a  rapid  stream  of  chlorine  takes  place,  thus  — 

K2Cr207  +  14HC1  =  2KC1  +  Cr2Cl6  +  7H2O  +  3C12. 

(4.)  When  crystals  of  potassium  chlorate  are  similarly  treated,  a 
mixture  of  chlorine  and  chlorine  peroxide  is  evolved,  even  without 
the  application  of  heat  — 

4KC1O3  +  12HC1  =  4KC1  +  6H2O  +  3C1O2  +  9C1. 

(5.)  Red  lead  (Pb3O4),  when  treated  with  hydrochloric  acid, 
reacts  in  a  manner  similar  to  manganese  dioxide  and  many  other 
peroxides.  In  the  case  of  lead,  however,  there  is  no  intermediate 
chloride  formed  — 

Pb3O4  +  8HC1  =  3PbCl2  +  4H2O  +  C12. 

(6.)  Manufacturing-  Processes  —  Deacon's  Process.  —  This 
method  for  the  preparation  of  chlorine  is  by  the  oxidation  of  the 
hydrogen  in  hydrochloric  acid  by  atmospheric  oxygen.  It  will  be 
seen  that  in  the  foregoing  methods  the  oxidation  of  this  hydrogen 
is  carried  on  at  the  expense  of  the  oxygen  contained  in  either  the 
metallic  peroxide  or  the  highly  oxygenated  salt  used  ;  in  the 
Deacon  process  atmospheric  oxygen  is  made  use  of.  When  a 
mixture  of  gaseous  hydrochloric  acid  and  oxygen  is  heated,  a 
slight  decomposition  takes  place  ;  but  if  these  gases  be  heated  in 
the  presence  of  a  third  substance  which  acts  as  a  catalytic  agent, 
the  decomposition  of  the  hydrochloric  acid  is  much  more  readily 
effected.  The  catalytic  agent  employed  in  the  Deacon  process  is 
cuprous  chloride  (Cu2Cl2).  This  substance  is  capable  of  taking  up 


Chlorine 


355 


an  additional  quantity  of  chlorine,   and  of  being  converted  into 
cupric  chloride  (CuCl2),  thus  — 


If,  therefore,  a  mixture  of  hydrochloric  acid  and  oxygen  be 
passed  over  fragments  of  pumice  impregnated  with  cuprous 
chloride  contained  in  a  tube  which  is  heated  to  dull  redness,  the 
hydrochloric  acid  will  be  decomposed.  We  may  suppose  that  the 
affinity  of  oxygen  for  the  hydrogen  in  hydrochloric  acid  is  un- 
able to  overcome  the  affinity  existing  between  the  hydrogen  and 
chlorine,  but  that  the  additional  pull  exerted  upon  the  molecules 
of  hydrochloric  acid  by  the  cuprous  chloride  is  sufficient  to  dis- 
turb the  equilibrium  and  rupture  the  molecule  — 


O 


H-C1 

H;CI' 


The  result  of  the  action  being  H2O  +  2CuCl2. 


FIG.  93. 

*  At  the  temperature  at  which  the  reaction  is  carried  on,  however, 
the  compound  CuCl2  cannot  exist  ;  two  molecules  of  it  are  con- 
verted into  one  of  Cu2Cl2,  and  a  molecule  of  chlorine  is  evolved. 
The  final  result,  therefore,  of  the  reaction  may  be  thus  expressed — 

O  +  2HC1  -f  Cu2Cl,  =  H2O  +  C12  4-  Cu2CI2. 

In  reality  the  action  is  rather  more  complex,  there  being  an  intermediate 
compound  formed  by  the  combination  of  cuprous  chloride  with  oxygen.     This 


356 


Inorganic  Chemistry 


oxychloride  of  copper  then  acts  upon  the  hydrochloric  acid,  as  seen  in  the 
following  equations  : — 

(i)  Cu2Cl2-fO=Cu2OCl2. 

(ft)  Cu2OCl2+2HCl-2CuCl2+H2O. 

(3)  2CuCl2=Cu2Cl2+Cl2. 

This  reaction  may  be  made  on  a  small  scale  by  means  of  the 
apparatus  shown  in  Fig.  93.  Hydrochloric  acid  is  generated  from 
salt  and  sulphuric  acid  in  the  flask,  and  a  stream  of  the  gas  passed 
through  the  Woulf's  bottle,  into  which  also  enters  a  stream  of 
oxygen.  The  mixed  gases  are  then  passed  through  the  bulb-tube, 
containing  fragments  of  pumice  which  have  previously  been  soaked 


FIG.  94. 

in  a  solution  of  cupric  chloride  and  dried.  On  heating  the  bulb  by 
means  of  a  Bunsen  flame,  chlorine  will  issue  from  the  end  of  the 
tube.  When  chlorine  is  manufactured  on  an  industrial  scale  by 
the  Deacon  process,  the  mixture  of  hydrochloric  acid  and  air  (in 
the  proportion  of  four  volumes  of  the  latter  to  one  volume  of  hydro-' 
chloric  acid)  is  drawn  by  means  of  a  Root's  blower  first  through 
iron  pipes,  which  are  heated  to  a  temperature  of  about  500°,  and 
then  the  hot  gases  pass  on  through  the  decomposer.  This  consists 
of  a  cylinder  of  cast  iron  containing  masses  of  broken  brick  or 
burnt  clay  impregnated  with  cupric  chloride,  and  so  arranged  that 
the  gases  are  drawn  through  the  masse 


Chlorine  357 

The  gas  leaving  the  decomposer  consists  of  a  mixture  of  chlorine, 
undecomposed  hydrochloric  acid,  and  atmospheric  nitrogen  and 
oxygen.  By  passing  them  through  water,  the  hydrochloric  acid  is 
removed,  and  the  chlorine  is  usually  converted  at  once  into  bleach- 
ing-powder. 

The  process  by  which  chlorine  is  usually  made  on  a  manufactur- 
ing scale  is  by  the  action  of  hydrochloric  acid  upon  manganese 
dioxide.  The  best  ore  for  the  purpose  is  pyrolusite.  The  process 
is  conducted  in  stills  made  of  thick  slabs  of  stone,  usually  "  York- 
shire flag,"  which  are  fitted  and  luted  together,  and  securely  bound 
by  cast-iron  clamps.  Fig.  94  shows  such  a  chlorine  still,  repre- 
sented as  cut  across  the  centre. 

The  charge  of  manganese  is  placed  upon  the  false  bottom  rt,  and 
the  acid  is  run  in  through  the  funnel  tube  £,  which,  dipping  into  a 
small  pot,  does  not  allow  the  gas  to  escape.  As  the  action  begins 
to  slacken,  steam  is  cautiously  blown  in  from  time  to  time.  The 
chlorine  escapes  by  the  pipe  g,  and  passes  from  thence  into  the 
main  h. 

The  reaction  that  goes  on  in  the  still  is  the  same  as  that  given 
above  in  the  first  mode  of  formation,  except  that  as  pyrolusite  is 
not  pure  MnO2,  small  quantities  of  other  compounds  are  formed. 
The  following  analysis,  by  Black,  of  still-liquor  from  stone  stills, 
shows  the  general  composition  of  this  substance  : — 

MnCl2 10.5700 

A12C16 0.6200 

Fe2Cl6 0.4551 

HC1  (undecomposed)     .         .         .  6.6220 

H2O 81.7329 

100.0000 

(7.)  The  Weldon  Process,  although  indirectly  a  method  for 
making  chlorine,  is  in  reality  a  process  for  recovering  the  man- 
ganese contained  in  the  still-liquors  as  manganous  chloride,  and  of 
reconverting  it  into  available  manganese  dioxide.  The  manganese 
so  recovered,  however,  is  again  utilised  for  the  preparation  of 
chlorine  by  the  decomposition  of  a  further  quantity  of  hydrochloric 
acid.  The  essence  of  the  process  is  the  following : — The  still- 
liquor  is  mixed  with  ground  chalk,  or  limestone  dust,  in  large  tanks 
or  wells,  and  the  mixture  thoroughly  stirred  by  agitators.  One  of 
these  wells,  A,  is  shown  in  the  diagrammatic  figure.  By  this  opera- 


358 


Inorganic  Chemistry 


tion  the  free  acid  is  neutralised,  and  the  iron  precipitated  as 
hydrated  oxide.  The  neutral  liquor,  consisting  of  manganous 
chloride  and  calcium  chloride,  is  then  pumped  into  large  tanks, 
where  it  is  allowed  to  settle ;  one  of  these  "  settlers,"  B,  is  shown 
in  the  figure.  By  means  of  a  pipe  upon  a  swivel-joint,  yj  the  clear 
liquid  from  the  settler  can  be  drawn  "off  without  disturbing  the 
sediment,  and  run  into  the  oxidiser  C.  The  oxidiser  is  merely  a 
flat-bottomed  iron  cylinder,  open  at  the  top.  Milk  of  lime  from 


FIG.  95. 

the  tank  E,  where  lime  and  water  are  stirred  together,  is  pumped 
into  the  oxidiser  as  required. 

The  milk  of  lime  is  added  in  quantity  more  than  sufficient  to 
precipitate  the  manganese  as  manganous  hydroxide,  MnH2O2. 
Into  this  mixture,  which  consists  of  manganous  hydroxide  and 
calcium  hydroxide  (milk  of  lime)  in  suspension,  and  to  a  smaller 
extent  in  solution  in  the  calcium  chloride  which  is  present,  a 
stream  of  compressed  air  is  forced  by  means  of  the  pipe  ^,  which 


Chlorine  359 

passes  to  the  bottom  of  the  oxidiser,  where  it  ends  in  perforated 
branches.  During  this  process  the  manganese  becomes  oxidised 
and  is  converted  mainly  into  calcium  manganite,  a  compound  of 
manganese  dioxide  with  calcium  oxide,  CaO,MnO2,  or  CaMnO3. 
By  a  further  addition  of  the  neutral  liquor  from  tank  B,  and  by 
raising  the  temperature  within  the  oxidiser  by  injecting  steam,  a 
portion  of  the  calcium  manganite  is  converted  into  a  compound 
having  the  composition  CaO,2MnO2. 

When  the  operation  is  complete,  the  contents  of  the  oxidiser  are 
run  out  into  a  series  of  tanks  called  mud  settlers,  of  which  one 
is  shown  at  D  in  the  figure.  The  product  here  settles  as  a  thin 
black  mud,  known  as  the  Weldon  mud;  and  this  is  ultimately 
drawn  from  the  settlers,  and  run  direct  into  chlorine  stills,  where 
it  is  at  once  treated  with  hydrochloric  acid  for  the  preparation 
of  chlorine.  The  Weldon  stills  are  similar  to  the  ordinary  chlorine 
stills,  but  are  much  larger,  and  usually  octagonal  in  shape. 

(8.)  Electrolytic  Methods.— Of  late  years,  since  the  applica- 
tion of  electricity  on  a  commercial  scale  has  become  possible, 
manufacturing  processes  for  obtaining  chlorine  by  the  electrolysis 
of  a  solution  of  common  salt  have  begun  to  compete  with  the 
older  methods.  By  the  electrolysis  of  brine,  the  sodium  chloride 
is  separated  into  its  two  elements  ;  the  chlorine  is  evolved  at  the 
anode,  and  the  sodium  which  is  liberated  at  the  cathode  there  acts 
upon  the  water  present,  generating  sodium  hydroxide  (see  Caustic 
Soda  ;  also  Sodium  Carbonate). 

Properties. — Chlorine  is  a  greenish-yellow  coloured  gas,  with  a 
strong  suffocating  smell.  It  is  quite  irrespirable,  and  if  inhaled  in 
the  pure  state  causes  death.  Even  when  largely  diluted  with  air 
it  is  extremely  disagreeable  and  injurious,  as  it  acts  rapidly  upon 
the  mucous  membranes  of  the  nose  and  throat,  causing  irritation 
and  inflammation,  which  usually  result  in  severe  catarrh.  A  few 
bubbles  of  chlorine  allowed  to  escape  and  diffuse  into  the  air  of  a 
room  give  to  the  air  a  distinct  and  rather  pleasant  smell.  Chlorine 

is  an  extremely  heavy  gas,  being    -      -  =  2.45  times  heavier  than 

14.44 
air.     One  litre  of  the  gas,  measured  under  the  standard  conditions 

of  temperature  and  pressure,  weighs  3.168  grammes.  The  density 
of  chlorine,  taken  at  all  temperatures,  does  not  exactly  agree  with 
that  which  is  required  for  the  molecular  formula  C12.  At  tempe- 
ratures above  1200°  the  density  is  markedly  less  than  theory 
demands,  showing  that  partial  dissociation  of  the  chlorine  mole- 


360  Inorganic  Chemistry 

cules  into  single  atoms  has  taken  place.  (Compare  Bromine  and 
Iodine.) 

On  account  of  its  heaviness,  chlorine  is  readily  collected  by  dis- 
placement ;  it  cannot  be  collected  over  mercury,  as  it  attacks  that 
metal,  and  in  water  it  is  considerably  soluble.  It  may,  however, 
be  collected  over  a  strong  brine,  as  it  is  much  less  soluble  in  this 
solution  than  in  water. 

Chlorine  is  not  inflammable,  but  it  supports  the  combustion  of 
many  burning  bodies.  It  is  possessed  of  such  extremely  powerful 
chemical  affinities  that  it  acts  upon  a  large  number  of  substances 
at  ordinary  temperatures,  and  in  many  cases  the  combination  is 
sufficiently  energetic  to  result  in  the  inflammation  of  the  bodies. 
Phosphorus,  when  introduced  into  chlorine,  first  melts  and  then 
spontaneously  inflames,  burning  with  a  somewhat  feeble  light  to 
form  phosphorus  trichloride  (PC13)  and  phosphorus  pentachloride 
(PC15).  The  elements  arsenic  and  antimony,  when  finely  powdered 
and  dusted  into  a  vessel  of  chlorine  at  once  take  fire  and  burn, 
forming  their  respective  chlorides.  Many  metals,  when  finely 
divided,  or  in  the  form  of  thin  leaf,  such  as  ordinary  Dutch  metal, 
instantly  take  fire  when  brought  into  chlorine.  If  a  quantity  o* 
sodium  be  heated  in  a  deflagrating  spoon  until  it  begins  to  burn  in 
the  air,  and  be  then  plunged  into  chlorine,  the  sodium  continues  to 
burn  in  the  gas  with  dazzling  brilliancy,  forming  sodium  chloride. 

Although  under  ordinary  circumstances  chlorine  unites  with 
metals  with  great  readiness,  it  has  been  shown  that  this  action  will 
not  take  place  if  the  chlorine  be  absolutely  dry.  Thus,  if  chlorine 
which  has  been  completely  freed  from  aqueous  vapour  be  passed 
into  a  tube  containing  bright  metallic  sodium,  and  the  tube  sealed 
the  sodium  not  only  remains  bright  and  unaffected  by  the  gas,  buf 
may  even  be  melted  in  the  atmosphere  of  chlorine  without  any 
action  taking  place.  Similarly,  dry  chlorine,  when  allowed  to  enter 
a  flask  filled  with  Dutch  metal,  has  no  action  upon  it  ;  but  upon  the 
introduction  of  the  smallest  trace  of  moisture  the  metal  at  once 
takes  fire.*  These  facts  are  of  the  same  order  as  those  mentioned 
in  connection  with  oxygen  (see  page  191). 

Chlorine  is  not  capable  of  direct  combination  with  carbon  ;  ordi- 
nary combustibles,  therefore,  which  consist  of  hydrocarbons,  burn 
in  chlorine  by  virtue  of  the  combination  of  their  hydrogen  with  the 
gas,  and  they  burn  with  a  lurid  smoky  flame,  owing  to  the  elimina- 
tion of  their  carbon  in  the  form  of  soot.  A  burning  taper  or 

*  See  Experiments  174,  175,  "  Chemical  Lecture  Experiments,"  new.  ed. 


Chlorine  361 

ordinary  gas  flame  when  introduced  into  chlorine  burns  in  this 
manner,  emitting  a  dense  smoke  and  forming  fumes  of  hydro- 
chloric acid. 

Chlorine  has  a  most  powerful  affinity  for  hydrogen  ;  a  jet  of 
hydrogen  burns  freely  in  chlorine,  with  the  formation  of  hydro- 
chloric acid.  A  mixture  of  hydrogen  and  chlorine  unites  with 
explosion  on  the  application  of  a  flame.  This  combination  takes 
place  also  under  the  influence  of  light  (see  Hydrochloric  Acid). 
The  affinity  shown  by  chlorine  for  hydrogen  is  seen  in  its  action 
upon  many  of  the  compounds  of  hydrogen  and  carbon.  If  one 
volume  of  ethylene  (C2H4)  be  mixed  with  two  volumes  of  chlorine, 
and  the  mixture  ignited,  the  carbon  is  instantly  thrown  out  of  com- 
bination as  a  black  smoke,  while  the  hydrogen  unites  with  the 
chlorine,  forming  a  cloud  of  hydrochloric  acid.  Similarly,  if  a 
liquid  hydrocarbon,  such  as  turpentine  (C10H10),  be  poured  upon  a 
piece  of  filter  paper,  and  the  paper  be  thrust  into  a  jar  of  chlorine, 
instant  inflammation  takes  place,  with  deposition  of  a  large  quantity 
of  carbon. 

Chlorine  possesses  strong  bleaching  properties,  which  depend 
upon  its  power  of  combining  with  hydrogen,  for  it  is  an  essential 
condition  that  water  shall  be  present.  The  chlorine  unites  with 
the  hydrogen  of  the  water,  and  the  liberated  oxygen  oxidises  the 
colouring  matter.  If  chlorine  be  bubbled  into  liquids  coloured 
with  any  vegetable  colouring  matter,  or  if  a  dyed  rag  be  dipped  into 
chlorine  water,  the  colour  will  be  rapidly  discharged.  Ordinary 
writing-ink  (which  usually  consists  of  a  compound  of  iron  with 
tannic  and  gallic  acids)  is  readily  bleached  by  chlorine  ;  while 
printer's  ink,  which  consists  mainly  of  carbon,  in  the  form  of  lamp- 
black, is  not  acted  upon  by  this  gas.  If,  therefore,  a  piece  of  printed 
paper  be  brushed  over  with  writing-ink  so  as  to  completely  obli- 
terate the  print,  and  the  blackened  paper  be  immersed  in  chlorine 
water,  the  writing-ink  will  be  rapidly  bleached  away,  leaving  the 
print  unchanged. 

The  bleaching  power  of  chlorine  constitutes  its  most  valuable 
property  from  an  industrial  point  of  view  ;  the  chlorine  for  this 
purpose  is  combined  with  lime  to  form  the  substance  known  as 
bleaching-powder  (see  Calcium  Compounds). 

Chlorine  is  soluble  in  water  to  a  considerable  extent.  One 
volume  of  water  at  10°  absorbs  3.0361  volumes  of  chlorine 
measured  at  o°  and  under  760  mm.  pressure.  This  solution, 
known  as  chlorine  water,  has  the  same  colour  as  the  gas,  and 


362  Inorganic  Chemistry 

smells  strongly  of  chlorine.  If  exposed  to  the  air,  the  chloi;ne 
rapidly  diffuses  out  of  the  solution.  Chlorine  water  cannot  be 
preserved  for  any  length  of  time,  as  it  slowly  undergoes  de- 
composition, the  chlorine  combining  with  the  hydrogen  of  the 
water,  forming  hydrochloric  acid,  which  remains  in  solution,  and 
the  oxygen  being  liberated,  thus — 

H2O  +  C12=2HC1  +  O. 

This  action,  which  proceeds  slowly  under  ordinary  conditions, 
is  greatly  accelerated  by  the  influence  of  light,  and  if  exposed 
to  direct  sunlight  the  decomposition  is  very  rapid. 

If  chlorine  water  be  cooled  to  within  one  or  two  degrees  of  the 
freezing-point  of  water,  or  if  chlorine  be  passed  into  ice-cold  water, 
a  solid  crystalline  compoirhd  of  chlorine  with  water  is  deposited. 
This  substance  is  termed  chlorine  hydrate,  and  has  a  composition 
expressed  by  the  formula  C12,10H2O.  The  compound  is  very  un- 
stable, and  when  exposed  to  the  air  it  melts  and  rapidly  gives  off 
chlorine.  If  the  crystals  are  quickly  freed  from  adhering  water, 
and  are  then  sealed  up  in  a  glass  tube,  they  may  be  heated  to 
a  temperature  of  38°  before  being  decomposed.  Faraday  made 
use  of  this  compound  in  order  to  obtain  liquefied  chlorine.  A 
quantity  of  the  hydrate  was  sealed  up  in  one  limb  of  a  bent 
tube  and  was  gently  warmed,  the  compound  dissociated  into 
water  and  chlorine,  and  the  internal  pressure  caused  the  condensa- 
tion of  the  chlorine  to  the  liquid  condition. 

Liquid  Chlorine. — Under  the  ordinary  atmospheric  pressure, 
chlorine  may  be  liquefied  by  lowering  its  temperature  to  —  34°. 

At  a  temperature  of  o°  the  pressure  required  to  effect  its  lique- 
faction is  equal  to  six  atmospheres.  When,  therefore,  the  liquid 
is  obtained  by  heating  the  crystalline  hydrate,  as  in  Faraday's 
method,  one  limb  of  the  tube  should  be  cooled  by  being  placed 
in  ice. 

The  critical  temperature  of  chlorine  is  141°,  and  the  pressure 
required  to  effect  its  liquefaction  at  that  point,  or  its  critical 
pressure,  is  84  atmospheres  (see  Liquefaction  of  Gases). 

Liquid  chlorine  has  a  bright  golden-yellow  colour,  entirely  free 
from  the  greenish  tint  possessed  by  the  gas.  Its  specific  gravity 
is  1.33,  and  it  boils  at  -33.6°.  When  cooled  to  a  temperature 
of  -  102°,  it  freezes  to  a  yellow  crystalline  mass.  Liquid  chlorine 
is  now  an  article  of  commerce.  It  is  contained  in  iron  bottles 
lined  with  lead,  and  is  largely  exported  in  this  form,  for  use  in  the 


Hydrochloric  Acid 


363 


extraction  of  gold,  to  parts  of  the  world  where  the  carriage  of  the 
plant  and  materials  necessary  for  generating  large  quantities  of 
chlorine  would  be  attended  with  great  difficulties. 


HYDROCHLORIC  ACID  (Hydrogen  Chloride). 
Formula,  HC1.     Molecular  weight  =  36. 46.     Density  =18.23. 

History.— In  solution  in  water  this  compound  was  known  to 
the  early  alchemists,  and  the  mixture  of  this  solution  with  nitric 
acid  constituted  the  valued  liquid  known  as  aqua  regia.  The 


FIG.  96. 

preparation  of  hydrochloric  acid  from  common  salt  is  associated 
with  the  name  of  Glauber  (1650),  who  obtained  it  by  the  action 
of  sulphuric  acid  upon  sodium  chloride  (common  salt).  Gaseous 
hydrochloric  acid  was  first  collected  and  examined  by  Priestley, 
who  collected  it  over  mercury  in  the  mercurial  pneumatic  trough 
invented  by  him.  He  named  the  gas  marine  acid  air. 

Occurrence.— Gaseous  hydrochloric  acid  is  evolved  in  consider- 
able quantities  from  volcanoes  during  active  eruption. 

Modes  Of  Formation.— (i.)  Hydrochloric  acid  may  be  syn- 
thetically produced  directly  from  its  elements  ;  thus,  this  compound 
is  formed  when  a  jet  of  hydrogen  is  caused  to  burn  in  an  atmos- 


364  Inorganic  Chemistry 

phere  of  chlorine.  If  a  mixture  of  chlorine  and  hydrogen  be 
ignited,  the  union  takes  place  instantaneously  with  explosion,  and 
hydrochloric  acid  is  produced.  The  union  of  hydrogen  with 
chlorine  will  also  take  place  under  the  influence  of  light  ;  thus,  if  a 
mixture  of  these  two  gases  be  exposed  to  even  diffused  daylight  for 
a  few  hours  the  greenish  colour  imparted  to  the  mixture  by  the 
chlorine  will  gradually  disappear,  and  on  examination  it  is  found 
that  the  tube  contains  hydrochloric  acid.  This  combination,  which 
is  only  gradual  when  the  mixture  is  exposed  to  diffused 
daylight,  becomes  explosively  sudden  if  the  mixed  gases 
are  exposed  to  direct  sunlight,  or  any  artificial  light  which 
is  rich  in  rays  of  high  refrangibility — the  so-called  actinic 
rays.  If,  therefore,  a  glass  vessel  be  filled  with  a  mixture 
of  these  gases  in  equal  volumes,  and  the  mixture  be  placed 
in  bright  sunshine,  a  violent  explosion  will  result,  and 
hydrochloric  acid  will  be  produced.  This  phenomenon  is 
best  illustrated  by  filling  small  thin  glass  bulbs  with  a 
mixture  of  the  two  gases  obtained  by  the  electrolysis  of 
aqueous  hydrochloric  acid.  The  bulbs  when  filled  can  be 
hermetically  sealed  before  the  blowpipe  without  causing 
the  combination  of  the  gases,*  and  if  kept  in  the  dark  may 
be  preserved  indefinitely. 

On  exposing  one  of  these  bulbs  to  the  light  of  burning 
magnesium  the  combination  of  the  two  gases  instantly 
takes  place,  with  a  sharp  explosion  which  shatters  the  bulb 
to  powder.  The  bulb  should  therefore  be  screened,  as 
shown  in  Fig.  96. 

The  rays  of  light  which  are  capable  of  causing  this 
combination  are  those  which  compose  the  blue  and  violet 
end  of  the  spectrum  ;  if  these  particular  rays  are  absorbed 
from  the  light  by  means  of  ruby  glass,  the  mixture  of 
FIG.  97.  gases  may  be  exposed  to  the  red  light  so  obtained  with- 
out any  action  taking  place.f 

The  combination  of  chlorine  with  hydrogen  is  not  attended  by 
any  alteration  in  volume  ;  one  volume  of  chlorine  combines  with 
onevolume  of  hydrogen,  and  the  resultant  hydrochloric  acidoccupies 
two  volumes.  This  may  be  readily  proved  by  filling  a  stout  glass 
tube,  provided  with  a  stop-cock  at  each  end,  with  a  mixture  of  the  two 
gases  in  exactly  equal  volumes,  and  causing  them  to  combine  either 

*  See  Experiment  178,  "Chemical  Lecture  Experiments,"  new  ed. 
t  Ibid.,  p.  93. 


Hydrochloric  Acid 


365 


by  the  influence  of  light  or  by  the  passage  of  an  electric  spark  by 
means  of  the  platinum  wires  sealed  into  the  tube  (Fig.  97).  On 
opening  one  of  the  stop-cocks  under  mercury  it  will  be  seen  that 
no  mercury  is  drawn  in,  neither  does  any  gas  pass  out  from  the 
tube,  thus  showing  that  the  union  has  taken  place  without  any 
alteration  in  the  volume.  If  one  of  the  cocks  be  now  opened 
beneath  water,  the  hydrochloric  acid  which  has  resulted  from  the 
union  of  the  hydrogen  and  chlorine,  being  extremely  soluble  in 
water,  the  liquid  will  rush  up  into  the  tube  and  completely  fill  it, 
showing  that  no  free  hydrogen  or  chlorine  remains  in  the  tube. 


FIG. 


(2.)  For  all  ordinary  purposes,  hydrochloric  acid  is  always 
obtained  by  the  action  of  sulphuric  acid  upon  sodium  chloride. 
For  laboratory  uses  the  apparatus  seen  in  Fig.  98  may  be  con- 
veniently employed.  Sulphuric  acid,  previously  diluted  with  rather 
less  than  its  own  volume  of  water,  is  placed  in  the  flask,  and  a 
quantity  of  common  salt  is  added.  On  the  application  of  a  gentle 
heat  a  steady  stream  of  gas  is  evolved,  which  may  be  dried  by 
being  passed  through  the  tubulated  bottle,  containing  pumice 
moistened  with  strong  sulphuric  acid.  The  gas  is  then  collected 
either  over  mercury  or  by  displacement.  The  reaction  which 
takes  place  is  expressed  by  the  equation — 

NaCl  +  H2SO4=  NaHSO4  +  HC1. 

If  strong  sulphuric  acid  be  employed  along  with  an  excess  of 
salt,  both  of  the  atoms  of  hydrogen  can  be  displaced  from  the 


366  Inorganic  Chemistry 

acid  ;  and  instead  of  the  hydrogen  sodium  sulphate  there  is 
formed  the  normal  sodium  sulphate — 

2NaCl  +  H2S04  =  Na2SO4  +  2HC1. 

A  much  higher  temperature  is  necessary  in  order  to  complete  the 
reaction  indicated  by  this  equation. 

Properties. — Hydrochloric  acid  is  a  colourless  gas  with  a 
choking,  pungent  odour.  In  contact  with  the  moist  air  it  forms 
dense  fumes,  consisting  of  minute  globules  of  a  solution  of  the 
gas  in  the  atmospheric  aqueous  vapour.  Hydrochloric  acid  does 
not  burn,  neither  does  it  support  ordinary  combustion. 

It  is  heavier  than  air,  its  specific  gravity  being — 

18.23          *'/••         \ 

J  =  1.26  (air=  i). 

14.44 

Hence  the  gas  is  readily  collected  by  displacement.  One  litre  of 
the  gas  weighs  18.185  criths. 

Hydrochloric  acid  is  extremely  soluble  in  water  ;  i  volume  of 
water  at  o°  and  under  a  pressure  of  760  mm.  is  capable  of  dis- 
solving 503  volumes  of  gaseous  hydrochloric  acid,  measured  at 
o°  and  760  mm.  As  the  temperature  rises  the  solubility  diminishes, 
as  seen  by  the  following  table  : — 

Temperature.  Coefficient  of  Absorption. 

0° 503 

30° 411 

50° 364 

The  solubility  of  hydrochloric  acid  may  be  illustrated  by  com- 
pletely filling  a  large  globular  flask  with  the  gas,  by  displacement, 
the  flask  being  provided  with  a  long  tube  passing  through  the  cork, 
as  seen  in  Fig.  99.  On  opening  this  tube  beneath  water,  the  gas 
begins  to  dissolve,  and  the  liquid  rises  slowly  in  the  tube  until  it 
reaches  the  top.  As  soon  as  the  first  few  drops  enter  the  globe, 
they  rapidly  absorb  the  gas,  thereby  causing  a  partial  vacuum  in 
the  vessel,  so  that  the  water  is  driven  up  the  tube  with  consider- 
able force,  forming  a  fountain,  which  continues  until  the  globe  is 
nearly  filled  with  liquid.  If  the  water  in  the  dish  is  rendered  blue 
by  the  addition  of  litmus  solution,  the  acid  nature  of  the  solution 
of  the  gas  will  be  evident  by  the  reddening  of  the  liquid  as"it 
enters  the  globe. 


Hydrochloric  Acid 


367 


When  a  weak  aqueous  solution  of  hydrochloric  acid  is  boiled 
it  loses  water  and  becomes  stronger  ;  while,  on  the  other  hand, 
if  a  strong  solution  be  heated,  it  loses  gas  and  becomes  weaker, 
until  in  both  cases  an  acid  containing  20.24  per  cent,  of  HC1  is 
produced  which  boils  at  110°.  This  strength  of  acid  corresponds 
to  a  composition  expressed  by  the  formula  HCl-f  8H2O,  and  it 
was  at  one  time  supposed  to  represent  a  definite  compound. 
Roscoe  and  Dittmar  have  shown,  however,  that,  as  with  nitric 
acid,  the  composition  of  the  liquid  which 
boils  at  a  constant  temperature  is  simply  a 
function  of  the  pressure.  (Compare  Nitric 
Acid,  page  239.) 

The  strongest  aqueous  solution  of  hydro- 
chloric acid  at  1 5°  C.  has  a  specific  gravity 
of  i. 212,  and  contains  42.9  per  cent,  of  HC1. 

Hydrochloric  acid  gas  is  readily  liquefied 
by  pressure.  At  a  temperature  of  10°  a 
pressure  of  40  atmospheres  will  effect  its 
liquefaction.  If  the  temperature  be  lowered 
to  —  1 6°,  the  same  result  is  obtained  by  a 
pressure  of  20  atmospheres.  The  critical 
temperature  of  hydrochloric  acid  is  52.3°. 

Condensed  hydrochloric  acid  is  a  colour- 
less liquid.  Gore  has  shown  that  this  lique- 
fied acid  is  without  action  on  most  of  the 


FIG.  99. 


metals  which  are  readily  dissolved  by  the  aqueous  acid. 

The  composition  of  hydrochloric  acid  may  be  experimentally 
proved  by  a  number  of  methods.  It  may  be  shown  synthetically 
by  the  volumetric  experiment  referred  to  above  (page  364). 

The  volumetric  proportion  of  hydrogen  contained  in  the  gas  may 
be  shown  by  means  of  sodium  amalgam.  The  sodium  in  the 
amalgam  acts  upon  the  hydrochloric  acid,  combining  with  the 
chlorine,  and  liberating  the  hydrogen  — 


For  this  purpose  gaseous  hydrochloric  acid  is  introduced  into  one 
limb  of  the  U-shaped  eudiometer  (Fig.  100),  and  its  volume  indicated 
by  means  of  a  ring  upon  the  tube,  the  mercury  being  level  in  both 
limbs.  A  second  ring  marks  exactly  half  the  volume.  A  quantity 
of  liquid  sodium  amalgam  is  then  poured  into  the  open  limb  until 
it  is  completely  filled,  and  on  being  closed  by  the  thumb  the  tube 


368 


Inorganic  Chemistry 


can  be  inverted  so  as  to  decant  the  gas  into  this  limb.  After  being 
bubbled  once  or  twice  through  the  amalgam,  the  gas  is  again 
returned  to  its  former  place  ;  and  by  drawing  mercury  from  the 
branch  tube,  the  levels  in  each  limb  can  be  again  adjusted,  when  it 
will  be  found  that  the  gas  remaining  in  the  tube  occupies  the  space 
exactly  down  to  the  upper  ring,  that  is  to  say,  two  volumes  of 
hydrochloric  acid  contain  one  volume  of  hydrogen.  That  the  gas 


FIG.  zoo. 


FIG.  101. 


FIG.  102. 


is  hydrogen  can  be  shown  by  again  filling  up  the  open  limb  with 
mercury,  and  driving  the  gas  out  of  the  stop-cock,  where  it  can  be 
inflamed  as  it  escapes. 

The  fact  that  hydrochloric  acid  contains  the  same  volume  of 
chlorine  as  of  hydrogen  may  also  be  demonstrated  by  collecting 
the  mixed  gases,  evolved  by  the  electrolysis  of  the  aqueous  acid,  in 
a  long  tube  provided  with  a  stoppered  funnel,  as  shown  in  Fig.  101. 


Hydrochloric  Acid  369 

The  gases  may  be  collected  over  a  saturated  solution  of  salt  in 
water,  and  the  tube  filled  to  the  lower  ring.  On  allowing  a  solution 
of  potassium  iodide  to  enter  by  means  of  the  funnel,  the  chlorine  is 
absorbed  with  the  liberation  of  iodine,  which  partially  dissolves 
and  partly  separates  as  a  solid.  When  the  absorption  of  the 
chlorine  is  complete,  the  water  will  have  risen  to  the  second  band 
placed  half-way  up  the  tube,  showing  that  one-half  of  the  gaseous 
mixture  consists  of  chlorine.  The  former  experiment  proved  that 
hydrochloric  acid  contained  half  its  volume  of  hydrogen,  therefore 
the  two  elements,  in  uniting  to  form  this  compound,  do  so  in  equal 
volumes  and  without  any  contraction  in  volume. 

When  aqueous  hydrochloric  acid  is  subjected  to  electrolysis,  the 
hydrochloric  acid  is  decomposed,  hydrogen  being  evolved  at  the 
negative  electrode  and  chlorine  at  the  positive.  At  first  the 
liberated  chlorine  is  dissolved  in  the  solution,  but  after  the  liquid 
has  become  saturated  with  the  gas,  the  whole  of  the  chlorine  is 
liberated.  By  conducting  this  decomposition  in  the  apparatus 
seen  in  Fig.  102,  and  continuing  the  passage  of  the  electric  current 
until  the  liquid  in  one  limb  is  saturated  with  chlorine  before  closing 
the  stop-cocks,  it  will  be  seen,  when  the  gases  are  collected  in  the 
tubes,  that  they  are  evolved  in  equal  volumes. 

The  Manufacture   of  Hydrochloric  Acid.  —  The  aqueous 

solution  of  hydrochloric  acid  is  an  object  of  commercial  manu- 

facture, which  is  carried  out  on  an  enormous  scale.     It  is  obtained 

>y  the  decomposition  of  common  salt  by  means  of  sulphuric  acid, 

according  to  the  reaction  — 


Formerly  hydrochloric  acid  was  a  waste  product  obtained  in  the  manufacture 
of  sodium  carbonate  by  the  method  known  as  the  Leblanc  process,  the  first 
stage  in  this  process  being  the  conversion  of  sodium  chloride  into  sodium 
sulphate  by  the  action  upon  it  of  sulphuric  acid.  The  hydrochloric  acid 
evolved  as  gas  in  this  reaction  was  allowed  to  escape  into  the  atmosphere. 
The  nuisance  caused  by  this  acid  gas  being  thrown  into  the  air,  ultimately 
resulted  in  the  "Alkali  Act,"  which  compelled  manufacturers  to  absorb  this 
waste  acid.  Since  that  time  the  Leblanc  process  for  the  manufacture  of  sodium 
carbonate  has  had  a  formidable  rival  in  another  method,  known  as  the 
ammonia-soda  process  (see  Sodium  Compounds),  which  would  probably  have 
completely  driven  the  older  method  out  of  the  field,  but  for  the  commercial 
value  of  the  hydrochloric  acid  which  is  obtained  as  a  secondary  product  in 
the  Leblanc  process.  The  hydrochloric  acid,  therefore,  which  formerly  was 
thrown  away  as  a  waste  product,  is  now  the  salvation  of  the  process,  and  the 
utmost  care  is  taken  to  prevent  any  of  it  from  escaping,  not  now  by  com- 
pulsion of  the  Alkali  Act  so  much  as  from  purely  economic  reasons. 

2  A 


Inorganic  Chemistry 

The  charge  of  salt  and  sulphuric  acid  is  heated  in  an  enormous 
hemispherical  cast-iron  pan,  built  into  a  brickwork  chamber,  so 
that  it  can  be  heated  by  a  fire  beneath,  and  so  that  the  evolved 
gas  can  be  conveyed  away  by  brick  or  earthenware  flues.  The  gas 
evolved  by  the  reaction  is  led  into  towers  which  are  filled  with  coke 
or  bricks,  and  down  which  water  is  made  to  percolate,  the  water 
being  caused  to  flow  equally  over  the  mass  by  means  of  special 
distributing  contrivances.  As  the  gaseous  hydrochloric  acid  passes 
up  the  towers  and  meets  the  descending  stream  of  water  it  is  en- 
tirely dissolved,  and  the  aqueous  acid  becomes  nearly  saturated  as 
it  reaches  the  bottom  of  the  tower. 

In  works  where   the   condensers   or  towers   are   not  of  great 


FIG.  103. 

height,  it  is  usual  either  to  cool  the  gas  before  admitting  it  into 
the  towers,  or  to  pass  it  through  a  series  of  jars  resembling  gigantic 
Woulf  s  bottles  (Fig.  103). 

The  water  in  these  bottles  is  made  to  flow  steadily  from  one  to 
the  other  by  the  side  pipes  c,  c  (in  the  direction  from  left  to  right), 
while  the  gas  passes  through  the  system  in  the  opposite  direction. 
In  this  way  a  constantly  changing  surface  of  water  is  exposed  to 
the  gas,  and  a  very  strong  solution  is  obtained. 

Commercial  hydrochloric  acid  is  generally  yellow  in  colour, 
owing  to  the  presence  of  iron  as  an  impurity  ;  and  it  is  always 
liable  to  contain  sulphuric  acid,  free  chlorine,  arsenic,  and  some- 


Chlorine  Monoxide  371 

times  sulphur  dioxide.  This  aqueous  solution  of  hydrochloric 
acid  is  also  known  under  the  names  of  "  spirits  of  salt "  and 
muriatic  acid. 


OXIDES   AND   OXY ACIDS   OF   CHLORINE. 

The  elements  oxygen  and  chlorine  have  never  been  made  to 
unite  together  directly  :  three  compounds,  however,  of  these  ele 
ments  can  be  obtained  by  indirect  methods  ;  these  are — 

Chlorine  monoxide  (hypochlorous  anhydride)        .     C12O. 
Chlorine  peroxide   .......     C1O2. 

Chlorine  heptoxide          ......     C12O7. 

Three  oxyacids  are  known,  viz.  : — 

Hypochlorous  acid HC1O. 

Chloric  acid HC1O3. 

Perchloric  acid HC1O4. 


CHLORINE  MONOXIDE  (Hypochlorous  Anhydride], 
Formula,  C12O.     Molecular  weight =86. 90.     Density =43. 45. 

Mode  Of  Formation. — This  compound  is  obtained  by  passing 
dry  chlorine  over  dry  precipitated  mercuric  oxide  contained  in  a 
glass  tube,  the  temperature  of  which  is  not  allowed  to  rise.  The 
chlorine  combines  with  the  mercuric  oxide,  forming  mercuric  oxy- 
chloride,  and  chlorine  monoxide  is  liberated — 

2HgO  +  2C12  =  HgO,HgCl2  +  C12O. 

Properties. — At  ordinary  temperatures  chlorine  monoxide  is  a 
pale  yellow  gas,  without  the  greenish  tint  possessed  by  chlorine. 
Its  smell  strongly  suggests  chlorine,  but  is  readily  distinguishable 
from  it.  It  is  a  very  unstable  compound,  decomposing  with  more 
or  less  violence  with  moderate  rise  of  temperature.  When  strongly 
cooled  it  is  condensed  to  an  orange-yellow  coloured  liquid,  which 
boils  at  about  —20°.  This  liquid  is  extremely  unstable,  exploding 
with  great  violence  on  the  gentlest  application  of  heat,  and  some- 
times on  merely  being  poured  from  one  vessel  to  another.  When 
exposed  to  direct  sunlight  it  also  explodes  with  violence. 

Gaseous  chlorine  monoxide  is  considerably  soluble  in  water,  one 


372  Inorganic  Chemistry 

volume  dissolving  about   100  volumes  of  the  gas,  forming  hypo- 
chlorous  acid — 

CLO  +  HoO  =  2HClO. 


CHLORINE  PEROXIDE. 

Formula,  C1O2.     Molecular  weight  =  67.  45.     Density  =33.  72. 

Modes  of  Formation.—  (  i.)  By  the  action  of  sulphuric  acid 
upon  potassium  chlorate  — 

3KC1O3  +  2H2SO4  =  KC1O4  :-  2H  KSO4  +  H2O  +  2C1O2. 

Finely  powdered  potassium  chlorate  is  added  little  by  little  to 
concentrated  sulphuric  acid  in  a  small  retort.  The  salt  dissolves 
with  the  formation  of  a  reddish  liquid,  and  if  the  temperature  is 
not  allowed  to  rise,  no  gas  is  evolved.  On  very  cautiously  warm- 
ing the  retort  by  means  of  warm  water,  taking'  care  not  to  heat 
the  glass  above  the  level  of  the  liquid  in  the  retort,  the  chlorine 
peroxide  is  evolved. 

(2.)  A  mixture  of  chlorine  peroxide  and  carbon  dioxide,  in  equal 
volumes,  is  obtained  by  heating  a  mixture  of  powdered  potassium 
chlorate  and  oxalic  acid  to  a  temperature  of  70°  in  a  water-bath  — 

2KC1O3  +  2H2C2O4=  K2C2O4  +  2H2O  +  2CO2  +  2C1O2. 

(3.)  Chlorine  peroxide,  mixed  with  chlorine,  is  evolved  by  the 
action  of  hydrochloric  acid  upon  potassium  chlorate  — 


This  mixture  of  gases  was  formerly  supposed  to  be  a  definite 
compound  of  oxygen  and  chlorine,  and  received  the  name  of 
euchlorine. 

Properties.  —  Chlorine  peroxide  is  a  heavy  gas,  with  a  deep 
yellow  colour.  It  has  an  intensely  unpleasant  smell,  and  if  in- 
haled, even  when  largely  diluted  with  air,  produces  headache. 
The  gas  attacks  mercury,  and  is  soluble  in  water,  so  that  it  can 
only  be  collected  by  displacement.  Chlorine  peroxide  is  an  ex- 
tremely unstable  compound,  it  is  gradually  resolved  into  its  ele- 
ments by  the  influence  of  light  ;  the  passage  of  an  electric  spark, 
or  the  introduction  into  it  of  a  hot  wire,  causes  it  to  decompose 
with  violent  explosion.  It  is  a  powerful  oxidising  compound  ;  a 


Hypochlorous  Acid  373 

piece  of  phosphorus  introduced  into  the  gas  takes  fire  spontane- 
ously. If  a  jet  of  sulphuretted  hydrogen  be  lowered  into  a  jar  of 
chlorine  peroxide,  the  sulphuretted  hydrogen  ignites  spontaneously 
and  continues  burning  in  the  gas. 

Its  oxidising  action  upon  organic  matter  may  be  shown  by 
liberating  the  gas  in  the  presence  of  such  a  substance  as  sugar, 
by  adding  a  drop  of  sulphuric  acid  to  a  mixture  of  powdered  sugar 
and  potassium  chlorate.  The  chlorine  peroxide,  liberated  by  the 
action  of  the  acid  upon  the  chlorate,  ignites  the  mixture,  and  the 
entire  mass  then  bursts  into  flame. 

When  chlorine  peroxide  is  strongly  cooled  it  condenses  to  a 
dark  red  liquid,  which  is  even  more  explosive  than  the  gas. 

Chlorine  HeptOXide,  C1.,O7.— This  compound  is  obtained  by 
the  cautiously  regulated  action  of  phosphoric  oxide  upon  perchloric 
acid,*  whereby  the  elements  of  water  are  withdrawn  from  two 
molecules  of  the  acid — • 

2HC1O4-H2O  =  C12O7. 

The  operation  is  attended  with  some  danger,  although  the  heptoxide 
when  isolated  is  described  as  less  unstable  than  either  of  the  other 
oxides. 

HYPOCHLOROUS  ACID. 

Formula,  HC1O. 

Modes  of  Formation.— ( i.)  As  already  mentioned,  this  acid  is 
formed  when  chlorine  monoxide  is  dissolved  in  water. 

(2.)  It  may  readily  be  obtained  in  dilute  solution  by  passing  an 
excess  of  chlorine  through  water  in  which  precipitated  mercuric 
oxide  is  suspended — 

HgO  4-  H2O  +  2C12=  HgCl2  +  2HC10. 

On  distilling  the  liquid,  the  dilute  acid  passes  over  as  a  colourless 
distillate. 

(3.)  In  dilute  solution,  hypochlorous  acid  may  be  obtained  by  the 
decomposition  of  a  hypochlorite  by  a  very  dilute  mineral  acid,  and 
subsequent  distillation  of  the  mixture  ;  thus,  if  to  a  solution  of 
calcium  hypochlorite  (obtained  by  treating  bleaching-powder  with 
water  and  filtering  the  solution)  very  dilute  nitric  acid  be  added 
and  the  solution  distilled,  a  dilute  colourless  acid  is  obtained — 

Ca(ClO)2  +  2HNO3  =  Ca(NO3)2  +  2HClO. 

(4.)  This  compound  is  also  formed  when  a  stream  of  chlorine  is 
*  Michael  and  Conn.,  Am.  Chem.  Journ.,  1900. 


374  Inorganic  Chemistry 

passed  through  water  containing  precipitated  calcium  carbonate  in 
suspension  — 

CaCO3+H2O  +  2Cl2  =  CaCl2  +  CO2  +  2HClO. 

Properties.—  Pure  hypochlorous  acid,  free  from  water,  has 
never  been  obtained.  The  acid  produced  by  the  solution  in  water 
of  chlorine  monoxide  has  a  pale  straw-yellow  colour,  and  a  very 
characteristic  chlorous  smell.  Dilute  solutions  of  this  acid  are 
moderately  stable,  while  more  concentrated  solutions  readily 
undergo  spontaneous  decomposition. 

Hypochlorous  acid  is  a  powerful  oxidising  and  bleaching  agent, 
as  it  readily  gives  up  its  oxygen,  and  is  resolved  into  hydrochloric 
acid  — 

HC1O  =  HC1  +  O. 

As  an  oxidising  agent  it  is  twice  as  effective  as  an  equivalent 
quantity  of  chlorine  in  chlorine  water,  for  two  atoms  of  chlorine  are 
here  necessary  for  the  liberation  of  one  atom  of  oxygen  — 


Hypochlorous  acid  is  decomposed  by  hydrochloric  acid   with 
the  evolution  of  chlorine  — 


It  is  also  decomposed  by  silver  oxide,  oxygen  being  liberated  — 


The  salts  of  hypochlorous  acid  may  be  obtained  by  the  action  of 
the  acid  upon  the  hydroxides  of  the  metals  ;  thus  — 


The  most  important  salt  of  this  acid  is  bleaching-powder  (see 
Calcium  Salts). 

CHLORIC    ACID. 

Formula,  HC1O3. 

Mode  Of  Formation.—  This  compound  is  best  obtained  by 
decomposing  barium  chlorate  with  an  exact  equivalent  of  sulphuric 
acid,  previously  diluted  with  water  — 

Ba(ClO3)2+H2SO4  =  BaSO4  +  2HClO3. 

The  clear  liquid  is  decanted  from  the  precipitated  barium  sul- 
phate, and  is  then  concentrated  by  evaporation  in  vacuo. 


Perchloric  Acid  375 

The  strongest  acid  that  can  be  obtained  still  contains  80  per 
cent,  of  water.  Attempts  to  concentrate  it  further  result  in  its 
decomposition  into  free  chlorine  and  oxygen,  with  the  formation 
of  perchloric  acid  and  water. 

Properties. — The  strong  aqueous  acid  has  powerful  oxidising 
properties  ;  many  organic  substances,  as  wood  or  paper,  are  so 
rapidly  oxidised  by  it  that  when  the  acid  is  dropped  upon  them 
they  are  frequently  inflamed. 

The  acid  even  in  dilute  solution  has  strong  bleaching  powers. 

The  salts  of  chloric  acid  are  far  more  stable  than  the  acid,  and 
some  of  them  are  of  considerable  technical  importance.  The 
chlorates  are  all  soluble  in  water,  and  all  yield  oxygen  on  being 
heated.  Chloric  acid  is  a  monobasic  acid  ;  the  chlorates,  there- 
fore, have  the  general  formula  M'C1O3  and  M"(C1O3)2,  where  M' 
and  M"  stand  for  monovalent  and  divalent  metals  respectively. 

Of  all  the  chlorates,  potassium  chlorate,  KC1O3,  is  by  far  the 
most  important  (see  Potassium  Compounds). 


PERCHLORIC  ACID. 

Formula,  HC1O4. 

Mode  of  Formation. — Perchloric  acid  is  best  prepared  by  the 
action  of  strong  sulphuric  acid  upon  potassium  perchlorate — 

2KC1O4+H2SO4  =  K2SO4  +  2HC1O4. 

Pure  and  dry  potassium  perchlorate  is  mixed  with  four  times  its 
weight  of  concentrated  sulphuric  acid,  and  the  mixture  gently  dis- 
tilled in  a  small  retort.  The  distillate  at  first  consists  of  perchloric 
acid  ;  but  as  the  operation  proceeds  a  portion  of  the  perchloric  acid 
is  decomposed  into  lower  oxides  of  chlorine  and  water,  and  the 
latter,  combining  with  the  first  portions  of  the  distillate,  forms  a 
white  crystalline  compound,  having  the  composition  HC1O4H2O. 
This  body,  when  gently  heated,  gives  off  perchloric  acid  ;  it  may, 
therefore,  be  employed  for  the  preparation  of  the  acid  in  a  state  of 
purity. 

Properties.— Perchloric  acid  is  a  colourless,  volatile,  and  strongly 
fuming  liquid,  having  a  specific  gravity  of  1.782  at  15°.  It  is  an 
extremely  powerful  oxidising  substance  ;  a  drop  of  the  liquid 
allowed  to  fall  upon  paper,  wood,  or  charcoal  is  instantly  decom- 
posed, sometimes  with  a  violent  explosion.  In  contact  with  the 


376  Inorganic  Chemistry 

skin  it  produces  most  painful  wounds  ;  when  allowed  to  drop  into 
water  it  produces  a  hissing  sound,  owing  to  the  energy  of  the 
combination. 

The  salts  of  this  acid  are  the  perchlorates,  of  which  the  most 
important  is  potassium  perchlorate  ;  they  are  all  soluble  in 
water. 

Constitution  of  the  Oxides  and  Oxyacids  of  Chlorine.— On  the  assump- 
tion that  chlorine  is  a  monovalent  element,  the  constitution  of  these  compounds 
may  be  thus  represented  : — 

Chlorine  monoxide,  Cl  -  O  -  Cl.     I    Hypochlorous  acid,  Cl  -  O  -  H. 

Chlorine  peroxide,  Cl  -  O  -  O  - .     I    Chloric  acid,  Cl-O-O-O-H. 

Perchloric  acid,  Cl-O-O-O-O-H. 

It  is  possible,  however,  that  in  some  of  these  compounds  the  chlorine 
functions  as  a  trivalent  element,  and  that  these  compounds  have  a  constitution 
similar  to  the  oxides  and  oxyacids  of  nitrogen,  thus  : — 


Chlorine  peroxide,  -  C1<T     |  .  Nitrogen  peroxide,  -  N 


Chlorine  monoxide,  Cl  -  O  -  Cl.  Nitrogen  monoxide,  N  —  O  -  N. 

/° 

Nitrogen  peroxide,  -  N\ 

^6 

Hypochlorous  acid,  Cl-  O  -  H.  Hyponitrous  acid,  N  -  O  -  H. 

/O  /O 

Chloric  acid,  H-O-Cl<f     |  .  Nitric  acid,  H  -  O  -  N^   |  . 

O 

Perchloric  acid,  H  -  O  -  Cl(      >O. 


O 

There  are  several  facts  which  point  to  the  belief  that  not  only  chlorine,  but 
also  bromine  and  iodine,  are  capable  of  fulfilling  the  functions  of  a  trivalent 
element.  The  existence,  for  example,  of  such  a  compound  as  trichloride  of 
iodine,  IC13,  is  difficult  to  explain  on  any  other  assumption  than  that  iodine  is 
here  a  trivalent  element. 

Indeed,  from  a  consideration  of  the  salts  of  periodic  acids,  some  chemists  are 
in  favour  of  assigning  to  iodine  even  a  still  higher  valency,  and  of  regarding  it 
as  a  heptad  element  in  these  compounds  (see  Periodates,  page  394).  The 
constitution  of  such  molecules  as  those  of  hydrofluoric  acid  at  low  temperatures, 
namely,  H2F2,  or  H3F3,  or  HnFn,  and  of  the  acid  fluoride  of  potassium, 
HF.KF,  is  readily  understood  if  we  regard  the  fluorine  as  functioning  in 
these  compounds  as  a  trivalent  element,  thus — 

H-F  =  F-H;  H-F  — F-H;  H-F=F-K. 

V 

H 


Bromine  377 


BROMINE. 

Symbol,  Br.     Atomic  weight  =  79.  92.     Molecular  weight  =159.  84. 
Vapour  density  =  79.  92. 

History.—  This  element  was  discovered  by  Balard  (1826),  in  the 
mother-liquor  obtained  after  the  crystallisation  of  salt  from  con- 
centrated sea-water.  He  applied  the  name  bromine  (signifying  a 
stencJi)  to  the  element,  in  allusion  to  its  unpleasant  smell. 

Occurrence.  —  Bromine  is  never  found  in  the  uncombined  state 
in  nature.  In  combination  chiefly  with  the  metals  potassium, 
sodium,  and  magnesium,  it  occurs  in  small  quantities  in  all  sea- 
water,  and  more  abundantly  in  many  mineral  waters  and  salt 
sprihgs.  The  saline  deposits  of  Stassfurt  contain  notable  quantities 
of  bromides,  and  the  main  supply  of  bromine  for  the  market  is 
manufactured  from  this  source. 

Modes  of  Formation.  —  (i.)  Bromine  may  be  obtained  from  a 
bromide  by  displacement  with  chlorine.  If  to  a  solution  of  mag- 
nesium bromide,  chlorine  water  is  added,  the  chlorine  combines 
with  the  magnesium  and  the  bromine  is  liberated  — 


On  distilling  the  liquid  the  bromine  is  driven  off,  and  can  be 
collected  in  a  well-cooled  receiver.  The  addition  of  any  excess  of 
chlorine  results  in  the  formation  of  bromide  of  chlorine,  and  is 
therefore  to  be  avoided. 

(2.)  Bromine  is  readily  obtained  from  potassium  bromide  by  the 
action  of  manganese  dioxide  and  sulphuric  acid,  a  reaction  exactly 
analogous  to  that  by  which  chlorine  is  obtained  from  sodium 
chloride  — 


The  mixture  is  gently  distilled  from  a  retort  into  a  receiver  kept 
cold  by  means  of  ice. 

(3.)  Manufacturing  Methods.—  Practically  all  the  bromine 
that  is  required  at  the  present  day  is  manufactured  from  crude 
carnallite  obtained  at  Stassfurt  (see  Alkali  Metals).  This  salt 
contains  bromine  combined  with  magnesium,  the  magnesium 
bromide  forming  about  i  per  cent,  of  the  magnesium  chloride  in 
the  crude  substance.  The  final  mother-liquors  from  the  manufac- 


378 


Inorganic  Chemistry 


ture  of  potassium  chloride,  and  which  were  formerly  run  to  waste, 
are  found  to  contain  about  .25  per  cent,  of  bromine  as  magnesium 
bromide,  and  these  liquors  are  now  utilised  for  the  manufacture  of 
bromine. 

The  bromine  is  liberated  from  its  combination  with  magnesium, 
by  means  of  chlorine.  In  some  processes  the  mother-liquor  is 
mixed  with  manganese  dioxide  and  sulphuric  acid  in  a  stone  vessel 


«-*  STEAM 


Cl 


FIG.  104. 

resembling  an  ordinary  chlorine  still.  The  magnesium  chloride  in 
the  liquor  is  acted  upon  by  the  manganese  dioxide  and  sulphuric 
acid  with  the  evolution  of  chlorine,  and  this  decomposes,  the 
bromide  present  displacing  the  bromine  — 


2  +  'Cl2=MgCl2 


The  bromine  that  is  driven  out  is  condensed  by  means  of  a  worm 
condenser. 


Bromine 


379 


Instead  of  the  chlorine  being  generated  within  the  mother-liquor, 
it  is  now  more  usually  produced  in  a  separate  chlorine  still,  and 
passed  into  the  liquor.  Fig.  104  shows  in  diagrammatic  form  the 
method  employed.  The  hot  mother-liquor  is  admitted  by  the  pipe 
A  into  the  tower  T,  which  is  filled  with  earthenware  balls,  between 
which  the  liquid  percolates.  It  leaves  the  tower  by  the  pipe  B, 
and  flows  into  the  tank  W,  which  is  provided  with  shelves  in  such 
a  way  that  the  liquid  must  circulate  through  it  in  the  direction 
indicated  by  the  arrows.  The  exit-pipe  from  this  tank  empties 
into  a  waste,  placed  at  such  a  height  that  the  tank  is  always  nearly 
full.  The  liquid  in  the  tank  is  kept  at,  or  near,  the  boiling-point,  by 
means  of  a  current  of  steam  blown  in  through  S.  Chlorine  from  a 


FIG.  105. 

still  is  admitted  by  the  pipe  L,  and  passing  into  the  tower  by  the 
pipe  B,  travels  in  an  opposite  direction  to  the  current  of  liquid. 
As  the  chlorine  passes  up  the  tower  it  meets  the  descending  mother- 
liquor,  and  decomposes  the  magnesium  bromide  contained  in  it 
with  the  liberation  of  bromine.  The  bromine  vapour  leaves  the 
cower  by  the  pipe  C,  and  is  conveyed  to  a  worm  (Fig.  105),  where  it 
is  condensed.  Any  bromine  which  dissolves  in  the  water  in  the 
tower  is  again  expelled  from  solution  by  the  steam  as  the  liquid 
traverses  the  tank  W,  and  is  swept  up  into  the  tower  by  the  current 
of  chlorine.  The  condensed  bromine  as  it  leaves  the  worm  is 
collected  in  a  tubulated  bottle,  and  any  vapour  which  escapes  con- 


380  Inorganic  Chemistry 

densation  is  arrested  by  the  vessel  F  (Fig.  105).  This  tube  is  filled 
with  irori  borings,  kept  moist  by  the  constant  dropping  of  water 
upon  them,  and  any  bromine  or  bromide  of  chlorine  is  there  con- 
verted into  iron  compounds,  which  are  dissolved  by  the  water, 
and  flow  away  into  the  receiver.  The  bromine  is  purified  by 
redistillation. 

Just  as  in  the  case  of  chlorine,  these  older  methods  of  manu- 
facture seem  destined  to  give  place  to  electrolytic  processes.  The 
method  now  being  introduced  for  the  extraction  of  the  bromine 
from  the  Stassfurt  liquors  depends  upon  the  fact  that  when  a 
solution  containing  a  chloride  and  a  bromide  is  submitted  to 
electrolysis,  the  bromine  is  liberated  first,  before  any  chlorine 
escapes.  Hence,  by  subjecting  the  liquors  to  electrolysis,  under 
suitable  conditions,  the  whole  of  the  bromine  is  readily  separated. 

Properties. — Bromine  is  a  heavy  but  mobile  liquid  of  a  deep 
reddish-brown  colour.  Except  in  extremely  thin  layers  it  is  opaque. 
It  is  the  only  non-metallic  element  which  is  liquid  at  the  ordinary 
temperature.  Bromine  boils  at  59°,  but  being  a  very  volatile  liquid 
it  gives  off  vapour  rapidly  at  the  ordinary  temperature.  A  drop  of 
bromine  allowed  to  fall  into  a  flask  immediately  evaporates  and 
fills  the  vessel  with  a  dark  red-brown  vapour.  The  specific  gravity 
of  the  liquid  at  o°  is  3.188.  At  -7°  bromine  solidifies  to  a  crystal- 
line mass.  Bromine  has  a  powerful  and  disagreeable  smell.  When 
the  vapour,  largely  diluted  with  air,  is  inhaled,  it  suggests  chlorine 
by  its  smell  and  by  its  action  upon  the  mucous  membrane  of  the 
throat  and  nose  ;  it  has  in  addition,  however,  a  most  irritating 
action  upon  the  eyes.  It  is  very  poisonous,  and  the  liquid  exerts  a 
corrosive  action  upon  the  skin  ;  it  produces  a  yellow  colour  when 
brought  in  contact  with  starch. 

The  vapour-density  of  bromine,  taken  at  moderately  high  tem- 
peratures, gradually  becomes  less  than  is  demanded  by  the  formula 
Br2,  showing  that  dissociation  takes  place.  In  the  case  of  bromine 
this  is  more  marked  than  with  chlorine. 

Bromine  is  soluble  in  water,  imparting  its  own  colour  to  the 
solution  which  is  known  as  bromine  water.  100  grammes  of  water 
at  o°  dissolve  3.60  grammes  of  bromine.  The  solubility  steadily 
diminishes  as  the  temperature  rises  :  at  20°  it  is  3.208,  and  at  30° 
it  is  3.126. 

When  bromine  water  is  cooled  to  o°  it  deposits  a  crystalline 
hydrate  similar  in  composition  to  the  hydrate  of  chlorine,  Br2,10H2O. 

Bromine  resembles  chlorine  in  its  chemical  attributes  ;  it  com- 


Hydrobromic  Acid  381 

bines  directly  with  metals  and  many  other  elements,  although  with 
less  energy  than  is  exhibited  by  chlorine.  A  fragment  of  arsenic, 
for  example,  when  dropped  upon  bromine,  ignites  and  burns  upon 
the  surface  of  the  liquid. 

Like  chlorine,  it  has  bleaching  properties,  due  to  its  power  of 
combining  with  hydrogen. 


HYDROBROMIC   ACID  (Hydrogen  Bromide}. 
Formula,  HBr.     Molecular  weight  =  80.  93.     Density=4o.46. 

Modes  of  Formation.  —  (i.)  Hydrobromic  acid  can  be  obtained 
by  the  direct  union  of  its  elements.  Bromine  vapour  and  hydrogen, 
when  mixed,  do  not  combine  under  the  influence  of  light  ;  neither 
does  such  a  mixture  explode  when  a  light  is  applied  to  it.  The 
mixture,  however,  may  be  caused  to  burn,  when  hydrobromic  acid 
is  formed  ;  or,  if  the  mixed  gases  be  passed  through  a  red-hot  tube, 
the  same  result  follows.  A  simple  method  of  preparing  hydro- 
bromic acid  synthetically  consists  in  passing  a  mixture  of  hydrogen 
and  bromine  vapour  over  a  spiral  of  platinum  wire,  maintained  at  a 
red  heat  by  means  of  an  electric  current.* 

(2.)  The  best  method  for  the  preparation  of  gaseous  hydrobromic 
acid  consists  in  dropping  bromine  upon  red  phosphorus  which  has 
been  moistened  with  a  small  quantity  of  water,  when  tribasic 
phosphoric  acid  is  formed,  and  hydrobromic  acid  is  liberated  — 

P  +  4H2O  +  5Br=H3PO4  +  5HBr. 

We  may  suppose  that  in  this  reaction  the  bromides  of  phosphorus 
are  formed  and  simultaneously  decomposed,  the  action  of  water 
upon  these  compounds  being  thus  expressed  — 


(3.)  Hydrobromic  acid  may  be  obtained  by  the  action  of  phos- 
phoric acid  upon  potassium  bromide  — 

3KBr+H3PO4  =  K3PO4  +  3HBr. 

(4.)  If  sulphuric  acid  be  employed  (as  in  the  formation  of  hydro- 
chloric acid  from  a  chloride),  free  bromine  is  simultaneously  pro- 

*  See  "Chemical  Lecture  Experiments,"  new  ed.,  No.  225. 


382  Inorganic  Chemistry 

duced,  owing  to  the  reduction  of  a  portion  of  the  sulphuric  acid 
by  the  hydrobromic  acid  which  is  first  evolved,  thus  — 

H2SO4  +  2HBr=  SO2  +  2H2O  +  Br2. 

(5.)  A  dilute  aqueous  solution  of  hydrobromic  acid  may  also  be 
obtained  by  passing  a  stream  of  sulphuretted  hydrogen  through 
bromine  water  — 


(6.)  Hydrobromic  acid  is  readily  obtained  by  the  action  of 
bromine  upon  certain  hydrocarbons,  such  as  turpentine  or  melted 
paraffin.  The  action  is  one  of  substitution,  one  atom  of  bromine 
replacing  one  atom  of  hydrogen  in  the  compound,  and  the  hydrogen 
so  displaced  combining  with  a  second  bromine  atom  to  form  hydro- 
bromic acid.  Thus,  if  the  hydrocarbon  be  represented  by  the 
general  formula,  CnH2n  +  2,  the  action  of  bromine  will  be  repre- 
sented thus  — 

CnH2n+2  +  Br2=CnH2n+1Br+HBr. 

Properties.  —  Hydrobromic  acid  is  a  colourless,  pungent-smell- 
ing gas,  which  fumes  strongly  in  the  air.  It  is  extremely  soluble 
in  water,  forming  an  acid  liquid  strongly  resembling  aqueous 
hydrochloric  acid. 

When  boiled,  this  solution  loses  either  acid  or  water,  until  it 
reaches  a  degree  of  concentration  at  which  it  contains  48  per  cent. 
of  hydrobromic  acid.  The  acid  of  this  strength  then  continues  to 
boil  unchanged  at  126°.  As  with  hydrochloric  acid,  the  strength 
of  the  liquid  which  boils  at  a  constant  temperature  depends  upon 
the  pressure. 

Hydrobromic  acid  is  decomposed  by  chlorine,  with  the  liberation 
of  bromine  — 


In  its  chemical  behaviour,  hydrobromic  acid  closely  resembles 
hydrochloric  acid,  and  this  resemblance  is  extended  to  the  bromides. 
All  bromides  are  soluble  in  water,  except  mercurous  bromide, 
silver  bromide,  and.  lead  bromide,  the  latter  being  slightly 
soluble, 


Hypobrovicus  Acid  383 

OXY  ACIDS    OF   BROMINE. 

No  oxides  of  bromine  corresponding  with  the  oxides  of  chlorine 
have  as  yet  been  obtained  ;  two  oxyacids,  however,  are  known,  viz.  :  — 

Hypobromous  acid          .         .         .     HBrO. 
Bromic  acid     .....     HBrO3. 

HYPOBROMOUS  ACID. 

Formula,  HBrO. 

Mode  Of  Formation.  —  An  aqueous  solution  of  hypobromous 
acid  may  be  obtained  by  shaking  together  a  mixture  of  bromine 
water  and  precipitated  mercuric  oxide,  the  reaction  being  ana- 
logous to  that  by  which  hypochlorous  acid  is  prepared  — 


Properties.  —  Hypobromous  acid  is  an  unstable  compound  ;  it 
breaks  up  on  distillation  into  oxygen  and  bromine.  By  heating  to 
40°  in  vacuo,  however,  it  can  be  distilled  without  decomposition. 
The  aqueous  liquid  so  obtained  has  a  pale-yellow  colour.  It 
readily  gives  up  its  oxygen,  and  is  a  strong  bleaching  agent  ; 
when  heated  to  about  60°  it  decomposes. 

Bromous  Acid,  HBrO2.  —  When  bromine  is  shaken  up  with  a  saturated 
solution  of  silver  nitrate,  the  resulting  liquid  is  believed  to  contain  bromous 
acid,  probably  produced  by  the  formation  first  of  hypobromous  acid,  and  its 
subsequent  oxidation  to  bromous  acid  *  — 

AgNO3  +  Br2+H2O=HBrO  +  AgBr+HNO3. 
:HBrO  +  OH2+2AgN03+Br2=HBr02+2HN03+2AgBr. 

The  acid  has  not  been  isolated,  nor  have  any  of  its  salts  been  obtained. 

BROMIC    ACID. 

Formula,  HBrO3. 

Modes  Of  Formation.—  (  i.)  This  acid  is  only  known  in  aqueous 
solution  ;  in  this  form  it  may  be  obtained  by  the  action  of  bromine 
upon  silver  bromate  in  the  presence  of  water  — 


The  insoluble  silver  bromide  separates  out,  and  the  aqueous 
acid  can  be  decanted  from  the  precipitate. 

(2.)  A  solution  of  this  acid,  mixed  with  hydrochloric  acid,  is  also 
formed  when  chlorine  is  passed  through  bromine  water  — 


Richards,  Jour,  Soc.  Chem.  Ind.t  1906. 


384  Inorganic  Chemistry 

(3.)  The  decomposition  of  barium  bromate  by  the  requisite 
weight  of  sulphuric  acid  affords  the  best  method  for  the  preparation 
of  a  pure  aqueous  solution  of  bromic  acid — 

Ba(BrO3)2  +  H2SO4=BaSO4  +  2HBrO3. 

Properties. — Bromic  acid  is  an  unstable,  strongly  acid  sub- 
stance, closely  resembling  chloric  acid.  The  aqueous  solution 
may  be  concentrated  in  vacuo  until  it  contains  about  50  per  cent, 
of  bromic  acid,  representing  a  composition  of  I  molecule  of  the 
acid  to  7  of  water.  Beyond  this  degree  of  concentration,  or  if 
heated  to  100°,  the  acid  decomposes  into  bromine,  oxygen,  and 
water. 

The  bromates  are  formed  by  reactions  similar  to  those  by  which 
the  chlorates  are  produced  ;  thus,  by  adding  bromine  to  a  solution 
of  potassium  hydroxide,  a  mixture  of  potassium  bromide  and 
bromate  is  obtained — 

6KHO  +  3Br2  =  5KBr+KBrO3  +  3H2O. 

A.nd  the  two  salts  can  be  separated  by  crystallisation,  owing  to  the 
greater  solubility  of  the  bromide. 

The  bromates  decompose  on  being  heated,  some  with  the 
liberation  of  oxygen  and  formation  of  bromide — 

KBrO3=KBr  +  3O, 

but  without  the  intermediate  production  of  a  perbromate.  Others 
give  off  their  bromine  as  well  as  a  part  of  the  oxygen  they  contain, 
leaving  the  metal  in  combination  with  oxygen — 

Mg(BrO3)2  =  MgO  +  Br2  +  5O. 


IODINE. 

Symbol,  I.     Atomic  weight  =126. 92.     Molecular  weight  =  253. 84. 
Vapour  density  =126. 92. 

History. — Iodine  was  discovered  by  Courtois  (1812),  who  ob- 
served that  a  beautiful  violet  vapour  was  evolved  during  his 
endeavours  to  prepare  nitre  from  liquors  obtained  by  lixiviating 
the  ashes  of  burnt  seaweed.  The  substance  was  subsequently 
investigated  by  Gay-Lussac. 


Iodine  385 

Occurrence.  —  Like  all  the  other  members  of  this  group  of  ele- 
ments, iodine  is  not  found  in  nature  in  the  uncombined  condition.* 
In  combination  it  occurs  associated  principally  with  potassium, 
sodium,  magnesium,  and  calcium,  as  iodides  and  iodates. 

Iodine  is  a  widely  distributed  element,  although  not  occurring 
in  more  than  small  quantities  in  any  particular  source.  Thus  it  is 
found  in  small  quantities  in  sea-water  and  in  both  marine  plants 
and  animals.  The  amount  of  iodine  in  seaweed  varies  with  diffe- 
rent plants  ;  generally  speaking,  those  from  greater  depths  contain 
more  than  weeds  which  grow  in  comparatively  shallow  waters. 


Dry  Weed, 

Drift  weed   \  Laminaria  digitata  (stem)         .     0.4535 


Laminaria  stenophylla      .         .     0.4777 

x-,  ,      (  Fucus  serratus  ....     0.0856 

Cut  weed      {   . 

(  Ascophyllum  nodosum     .         .     0.0572 

Iodine  is  also  found  in  small  quantities  in  many  mineral  waters 
and  medicinal  springs. 

In  small  quantities  iodine  is  present  in  the  natural  sodium  nitrate 
of  Chili  and  Peru,  known  as  Chili  saltpetre,  and  at  the  present  day 
this  constitutes  the  most  abundant  source  of  this  element. 

Mode  of  Formation.  —  Iodine  may  be  readily  obtained  by  a 
precisely  similar  reaction  to  that  by  which  both  bromine  and 
chlorine  are  produced  ;  thus,  if  potassium  iodide  be  mixed  with 
manganese  dioxide  and  sulphuric  acid,  and  the  mixture  gently 
heated  in  a  retort,  iodine  distils  over  and  condenses  in  the  form  of 
greyish-black  crystals  — 

2KI  +  MnO2  +  2H2SO4  =  K2SO4+MnSO4+2H2O  +  Io. 

Manufacturing"  Processes.  —  On  an  industrial  scale  iodine  is 
obtained  from  two  sources,  namely,  from  seaweed  and  from  caliche 
(Chili  saltpetre). 

(i.)  From  seaweed.  The  weeds  chiefly  employed  are  the  Lami- 
naria digitata  and  Laminaria  stenophylla.  The  weed  is  burnt  in 
shallow  pits,  care  being  taken  to  avoid  too  high  a  temperature  ;  the 
maximum  yield  of  iodine  being  obtained  if  the  ash  is  not  allowed 
to  fuse.  This  ash  is  technically  known  as  kelp^  and  if  the  weed  is 
properly  burnt,  it  should  yield  a  kelp  containing  from  25  to  30  Ibs. 
of  iodine  per  ton.  The  kelpers,  however,  usually  lose  about  half 

*  It  is  on  record  (Wanklyn,  Chem.  News,  54)  that  a  minute  quantity  of  free 
iodine  was  found  i:i  the  water  from  Woodhall  Spa. 

2   B 


386  Inorganic  Chemistry 

the  iodine  on  account  of  burning  the  weed  at  too  high  a  tempera- 
ture, thereby  fusing  the  ash  into  a  hard  slag,  instead  of  obtaining 
a  porous  residue. 

An  improved  process  of  carbonising  the  weed  was  introduced  by 
Stanford  (1863),  in  which  it  was  heated  in  large  retorts,  whereby 
the  volatile  products  of  the  distillation,  consisting  largely  of  tar 
and  ammoniacal  liquor,  could  be  .collected.  The  kelp  obtained 
by  this  method  is  in  a  very  porous  condition,  and  contains  the 
whole  of  the  iodine  originally  present  in  the  weed. 

A  still  more  recent  process  for  extracting  the  iodine  from  sea- 
weed, and  at  the  same  time  obtaining  other  useful  materials,  has 
since  been  discovered  by  Stanford.  The  weed  is  boiled  with 
sodium  carbonate  and  filtered  :  the  residue  consists  of  a  substance 
called  algulose.  Hydrochloric  acid  is  added  to  the  filtered  liquid, 
which  precipitates  a  compound  known  as  alginic  arid,  and  this 
is  again  separated  by  filtration.  The  liquor  is  neutralised  with 
sodium  hydroxide,  evaporated  to  dryness  and  carbonised.  The 
residue,  which  is  known  as  "kelp  substitute,"  contains  all  the 
iodine,  as  well  as  the  potash  salts,  and  should  yield  about  30  Ibs. 
of  iodine  per  ton. 

[The  alginic  acid  obtained  in  this  process  is  purified  and  converted  into  the 
sodium  salt,  which  constitutes  the  commercial  "algin,"  a  material  of  a  gelatin- 
ous or  albuminous  nature  which  has  lately  been  put  to  a  number  of  useful 
applications.] 

The  kelp  obtained  by  either  of  these  methods  is  lixiviated  witr 
water  in  large  iron  vats,  whereby  all  the  soluble  salts  are  extracted. 
This  aqueous  liquid  is  concentrated  in  large  open  boiling  pans, 
and  the  less  soluble  salts,  viz.,  the  alkaline  sulphates,  carbonates, 
and  chlorides,  are  allowed  to  crystallise.  The  mother-liquor  is 
then  mixed  with  sulphuric  acid  and  allowed  to  stand.  The  sul- 
phuric acid  decomposes  any  sulphides  and  sulphites  which  may 
be  present,  with  the  separation  of  sulphur  ;  it  also  converts  the 
bromides  and  iodides  into  the  corresponding  sulphates,  with  the 
liberation  of  hydrobromic  and  hydriodic  acids  which  remain  in 
solution,  while  the  alkaline  sulphates  are  deposited  from  the  liquid, 
and  are  technically  known  as  plate  sulphate.  The  liquor  is  then 
transferred  to  the  iodine  still,  which  is  an  iron  pot  furnished  with 
a  leaden  cover  into  which  two  exit-pipes  are  fixed  (Fig.  106). 
These  are  connected  to  a  series  (usually  ten  in  each  row)  of  large 
earthenware  jars  or  aludels.  A  gentle  heat  is  applied,  and 


Iodine  387 

manganese  dioxide  is  introduced  from  time  to  time  through 
the  opening.  The  iodine  is  evolved  according  to  the  following 
equation  — 


2HI  +  MnO2+H2SO4=MnSO4 

and  condenses  in  the  jars.  These  vessels  are  also  furnished  with  a 
tubulus  upon  their  under  side,  so  that  the  water  which  is  evolved 
during  the  distillation  can  drain  out,  and  run  off  down  the  trough  in 
which  the  jars  are  resting. 

(2.)  From  Chili  saltpetre.     The  crude  sodium  nitrate  of  Chili 


FIG.  106. 

and  Peru,  known  as  caliche,  contains  small  quantities  of  iodine, 
chiefly  as  sodium  iodate.  Although  the  amount  of  iodine  in 
caliche  is  only  very  small,  averaging  about  0.2  per  cent.,  in  view 
of  the  enormous  quantity  of  nitrate  that  is  turned  out,  the  aggre- 
gate amount  of  iodine  is  very  great.  The  iodine  is  now  extracted, 
and  the  supply  of  this  element  that  is  now  manufactured  from  this 
source  is  more  than  the  total  consumption  of  iodine  in  the  whole 
world.  The  process  is  based  upon  the  fact,  that  when  a  solution 
of  hydrogen  sodium  sulphite  (sodium  bisulphite)  is  added  to  a 
solution  of  an  iodate,  iodine  is  precipitated,  thus  — 


The  final  mother-liquor  from  the  sodium  nitrate,  or  caliche,  in 
which  all  the  iodate  has  concentrated,  contains  as  much  as  22 
per  cent,  of  this  salt.  This  liquor  is  mixed  with  the  requisite 
proportion  of  the  hydrogen  sodium  sulphite  solution,  in  large  lead- 


388  Inorganic  Chemistry 

lined  vats,  and  the  precipitated  iodine  allowed  to  settle.  It  is 
then  washed  and  pressed  into  blocks,  and  is  found  to  contain 
from  80  to  85  per  cent,  of  iodine.  This  impure  product  is  then 
distilled  at  a  gentle  heat  from  iron  retorts,  the  vapour  being  con- 
densed in  a  series  of  earthenware  receivers  much  as  in  the  older 
method. 

Properties. — Iodine  is  a  bluish-black  shining  solid,  somewhat 
resembling  graphite  in  lustre  and  general  outward  appearance.  It 
crystallises  in  large  brilliant  plates,  which  have  a  specific  gravity 
of  4.95.  When  heated  to  107°  iodine  melts  and  gives  off  vapour 
having  a  beautiful  violet  colour.  Its  boiling-point  is  about  175°. 

Iodine  vaporises  slowly  at  ordinary  temperatures,  and  sublimes 
from  one  part  to  another  of  a  bottle  in  which  a  small  quantity  of 
it  is  contained.  The  smell  of  iodine  vapour  is  somewhat  irritating 
and  unpleasant,  recalling  the  smell  of  moderately  diluted  chlorine. 
When  iodine  vapour  is  heated  it  passes  from  a  violet  colour  to  a 
deep  indigo  blue.*  This  change  in  the  colour  is  accompanied 
by  a  diminution  of  the  vapour-density.  Up  to  a  temperature  of 
700°  the  density  of  iodine  corresponds  to  the  formula  I2 ;  as  the 
temperature  is  raised  the  density  gradually  diminishes,  until  at 
1468°  it  is  reduced  to  less  than  two-thirds.  At  this  point,  73.1  per 
cent,  of  the  iodine  molecules  have  become  dissociated  into  single 
atoms. 

Iodine  is  slightly  soluble  in  water,  I  gramme  of  iodine  being 
dissolved  by  5.524  litres  of  water  at  10°.  This  dilute  solution, 
however,  has  a  perceptible  brown  colour.  Iodine  is  freely  soluble 
in  aqueous  potassium  iodide  solution,  in  alcohol,  ether,  and  aqueous 
hydriodic  acid  ;  in  all  these  solvents  it  dissolves  to  a  dark  reddish- 
brown  solution.  In  chloroform,  carbon  disulphide,  and  many 
liquid  hydrocarbons,  iodine  is  also  soluble,  but  in  these  solvents 
it  dissolves  to  a  deep  violet  solution  resembling  the  colour  of  the 
vapour. 

When  iodine  is  brought  into  contact  with  starch  it  forms  an 
intense  blue  colour.  This  reaction  is  so  extremely  delicate  that 
it  is  capable  of  revealing  the  minutest  trace  of  iodine.  The  exact 
nature  of  this  blue  compound  is  not  known.  The  colour  disappears 
when  the  liquid  is  heated  to  about  80°,  but  returns  on  cooling  ; 
continued  boiling  destroys  it  permanently. 

In  its  chemical  relations  iodine  resembles  chlorine  and  bromine, 
but  with  a  lesser  degree  of  energy.  Both  these  elements  are 
*  "  Chemical  Lecture  Experiments,"  new  ed. ,  No.  231. 


Hydriodic  Acid 


389 


capable  of  displacing  iodine  from  its  combinations  with  electro- 
positive elements,  thus — 

KI  +  Br=KBr+I. 
KI+C1-KCL+I. 

Iodine  combines  with  many  elements,  both  metals  and  non- 
metals,  forming  iodides.  Phosphorus,  when  brought  in  contact 
with  iodine,  at  once  melts  and  inflames  ;  antimony  powder  dropped 
into  iodine  vapour  also  spontaneously  inflames.  When  mercury 
and  iodine  are  gently  heated,  energetic  combination  takes  place, 
and  mercuric  iodide  is  formed. 


HYDRIODIC  ACID  (Hydrogen  Iodide]. 
Formula  HI.     Molecular  weight  — 127.93.     Density  =  63. 96. 
Modes  Of  Formation.— ( i.)  Hydriodic  acid  can  be  obtained 


FIG.  107. 

synthetically  by  passing  a  mixture  of  hydrogen  and  iodine  vapour 
over  strongly  heated,  finely  divided  platinum. 

(2.)  It  is  also  obtained  by  the  action  of  phosphoric  acid  upon 
sodium  or  potassium  iodide  (see  Hydrobromic  Acid). 

As  in  the  case  of  the  corresponding  bromine  compound,  sul- 


3QO  Inorganic  Chemistry 

phuric  acid  cannot  be  employed,  as  by  its  action  upon  the  iodide, 
iodine  and  sulphur  dioxide  are  liberated,  thus  — 


(3.)  Hydriodic  acid  is  produced  by  the  action  of  sulphuretted 
hydrogen  upon  iodine  (p.  410).  At  the  ordinary  temperature, 
and  in  the  absence  of  water,  these  two  substances  do  not  react, 
hydriodic  acid  being  an  endothermic  compound  (p.  168)  ;  but  if 
the  iodine  be  suspended  in  water  and  sulphuretted  hydrogen 
passed  through,  the  heat  of  solution  of  the  hydriodic  acid  supplies 
the  necessary  energy  to  enable  the  action  to  proceed.  When, 
however,  the  solution  reaches  a  sp.  gr.  of  1.56  the  action  ceases, 
because,  as  Naumann  has  shown,  the  heat  produced  by  the  solution 
of  the  product  is  insufficient  to  carry  on  the  process  beyond  this 
degree  of  concentration. 

(4.)  Hydriodic  acid  is  most  readily  prepared  by  the  action  of 
phosphorus  upon  iodine  in  the  presence  of  water  — 


The  red  phosphorus  and  iodine  for  this  reaction  may  be  placed 
in  a  dry  flask,  and  water  gradually  dropped  upon  the  mixture,  when 
hydriodic  acid  is  rapidly  evolved.  The  gas  is  allowed  to  pass 
through  a  U-tube  containing  red  phosphorus,  in  order  to  arrest  any 
iodine  vapour  which  may  accompany  it.  Hydriodic  acid  may  be 
collected  over  mercury  or  by  displacement,  as  shown  in  Fig.  107. 

Properties.  —  Hydriodic  acid  is  a  colourless,  pungent-smelling 
gas,  which  fumes  strongly  on  coming  into  the  air.  The  gas  is 
readily  decomposed  by  heat  into  hydrogen  and  iodine.  Thus,  if  a 
haated  wire  be  thrust  into  the  gas,  or  if  a  spiral  of  platinum  wire 
be  heated  in  the  gas  by  means  of  an  electric  current,  the  violet 
vapour  of  iodine  at  once  makes  its  appearance. 

When  mixed  with  chlorine,  hydriodic  acid  is  at  once  decomposed, 
with  the  liberation  of  iodine,  thus  — 


Hydriodic  acid  is  one  of  the  most  readily  liquefied  gases  ;  at  o°, 
and  under  a  pressure  of  four  atmospheres,  it  condenses  to  a  colour- 
less liquid. 

The  gas  is  extremely  soluble  in  water.  An  aqueous  solution  of 
it  is  readily  produced  by  allowing  the  gas,  obtained  by  the  method 


Iodine  Pentoxide 


391 


of  preparation  above  described,  to  pass  into  water.     In  order  to 

prevent  the  water  from  being  drawn  back  into  the  generating  flask, 

it  is  convenient  to  pass  the  gas  through  a  retort  arranged  in  the 

position  seen  in  P'ig.  108.    Should  there  be 

any  back  rush  of  water,  owing  to  the  inter- 

mission of  the  evolution  of  gas  in  the  ap- 

paratus, the  liquid  in  the  beaker  will  be 

drawn  up  into  the  retort  and  there  lodge, 

leaving  the  end  of  the  neck  open  to  the 

air. 

A  saturated  aqueous  solution  of  hydri- 
odic  acid  at  o°  has  a  specific  gravity 
of  2.  At  the  ordinary  pressure  the 
strongest  acid  that  can  be  obtained  by 
distillation  has  a  specific  gravity  of  1.67, 
and  contains  57.7  per  cent,  of  hydriodic 
acid.  This  solution  boils  at  127°.  As 
in  the  case  of  the  corresponding  bromine 
and  chlorine  compounds,  the  particular 

strength  of  acid  which  has  a  constant  boiling-point  is  a  function 
of  the  pressure. 

Aqueous  hydriodic  acid,  when  freshly  prepared,  is  colourless,  but 
it  rapidly  turns  brown,  owing  to  the  oxidation  of  the  compound 
and  the  solution  of  the  liberated  iodine  in  the  acid  — 


FIG.  108. 


OXIDE  AND   OXY  ACIDS   OF  IODINE. 

One  *  compound  of  iodine  with  oxygen  is  well  known,  and  three 
oxyacids,  viz.  :  — 

Iodine  pentoxide  ....     IgOs- 
lodicacid     .....     HIO3. 
Periodic  acid         .         .         .         .     HIO4. 
Hypoiodous  acid  .         .         .     HIO. 

IODINE    PENTOXIDE    (lodic  Anhydride}. 
Fcrmula,  loOs. 

Mode  of  Formation.  —  When  ioclic  acid  is  heated  to  170°,  it 
loses  water  and  is  converted  into  the  pentoxide  — 
2HIO3=H2O+IaO6. 

*  A  second  oxide,  ICX  or  I.2O4,  has  recently  been  described  by  Muir,  Jour. 
Chem.  Soc.,  April  1909. 


392  Inorganic  Chemistry 

Properties. — Iodine  pentoxide  is  a  white  crystalline  solid  body. 
It  is  soluble  in  water,  combining  with  a  molecule  of  the  water  to 
form  iodic  acid.  Iodine  pentoxide  is  more  stable  than  any  of  the 
oxides  of  the  other  halogens,  but,  when  heated  to  a  temperature  of 
300°,  it  decomposes  into  its  elements. 


IODIC    ACID. 

Formula,  HIO3. 

Modes  Of  Formation.—  (  i.)  Iodic  acid  can  be  prepared  by 
adding  to  a  solution  of  barium  iodate  the  requisite  amount  of 
sulphuric  acid  demanded  by  the  equation  — 

Ba(I03)2+H2S04=BaS04  +  2HI03. 

The  aqueous  solution  of  iodic  acid  is  decanted  from  the  preci- 
pitated barium  sulphate,  and  may  be  concentrated  at  100°  without 
being  decomposed. 

(2.)  When  chlorine  is  passed  through  water  in  which  powdered 
iodine  is  suspended,  a  mixture  of  iodic  acid  and  hydrochloric  acid 
is  produced  — 

3H2O  +  I+5C1  =  5HC1  +  HIO3. 

The  hydrochloric  acid  may  be  removed  by  the  addition  of  preci- 
pitated silver  oxide  to  the  solution,  and  separating  the  precipitated 
silver  chloride  by  filtration. 

(3.)  Iodic  acid  is  most  conveniently  prepared  by  heating  iodine 
with  nitric  acid,  whereby  the  iodine  is  oxidised  and  a  mixture  of 
oxides  of  nitrogen  is  evolved  as  dense  'red  vapours  — 


Properties.  —  Iodic  acid  is  a  white  crystalline  solid,  soluble  in 
water.  The  aqueous  solution  shows  an  acid  reaction  with  litmus, 
but  the  colour  is  ultimately  discharged  by  the  bleaching  action  of 
the  compound.  Iodic  acid  does  not  form  any  blue  colour  with 
starch  ;  being,  however,  an  oxidising  substance,  it  readily  gives  up 
oxygen  to  such  reducing  agents  as  sulphur  dioxide,  sulphuretted 
hydrogen,  or  hydriodic  acid,  with  the  liberation  of  iodine,  thus  — 


+  5HI=3H 


Periodic  Acid  393 

If,  therefore,  a  small  quantity  of  sulphurous  acid  be  added  to  a 
dilute  solution  of  iodic  acid,  previously  mixed  with  starch,  the  blue 
iodide  of  starch  will  be  formed.  This  reaction  affords  an  excellent 
illustration  of  the  time  required  for  certain  chemical  changes  to  go 
forward.  It  is  readily  possible  to  obtain  an  interval  of  30  to  60 
seconds  between  the  addition  of  the  sulphurous  acid  and  appear- 
ance of  any  visible  result,  when  at  the  expiration  of  that  time  the 
whole  mass  of  the  liquid  suddenly  turns  blue.* 

lodates.  —  When  iodine  is  dissolved  in  potassium  hydroxide,  a 
mixture  of  potassium  iodide  and  iodate  is  produced  by  an  analogous 
reaction  to  that  which  takes  place  with  either  bromine  or  chlorine  — 


With  the  exception  of  the  iodates  of  the  alkali  metals,  the  iodates 
are  for  the  most  part  insoluble  in  water.  On  being  heated  they 
behave  in  a  similar  manner  to  the  bromates,  some  being  decom- 
posed into  an  iodide  and  oxygen,  while  others  leave  a  metallic 
oxide  and  evolve  iodine  as  well  as  oxygen.  The  alkaline  iodates 
are  capable  of  uniting  with  iodic  acid,  forming  salts  which  are 
termed  acid  and  di-acid  iodates,  thus  — 

Normal  potassium  iodate     .         .         .     KIO3. 
Acid  potassium  iodate          .         .         .     KIO3,HIO3. 
Di-acid  potassium  iodate     .         .        .     KIO3,2HIO3. 


PERIODIC  ACID. 

Formula,  HIO4,2H2O  or  H5IO6. 

Modes  of  Formation.—  (  i.)  The  compound  represented  by  the 
formula  HIO4has  never  been  obtained;  when  aqueous  solutions 
of  periodic  acid  are  evaporated,  the  compound  which  crystallises 
out  has  the  composition  HIO4,2H2O,  or  H6IO6. 

It  may  be  obtained  by  boiling  silver  periodate  with  water,  when 
an  insoluble  basic  silver  salt  is  produced  — 


=  Ag2H3I06+HI04,2H20. 

The  silver  periodate  is  prepared  by  passing  chlorine  into  an  aqueous  solu- 

*  See  Experiment  246.  "Chemical  Lecture  Experiments,"  new  ed. 


394  Inorganic  Chemistry 

tion  of  sodium  iodate  and  sodium  hydroxide,  when   the   sparingly  soluble 
disodium  periodate  separates  out  — 

NaIO3  +  3NaHO  +  C12  =  2NaCl  +  Na2H3IO6. 

This  sodium  salt  is  then  dissolved  in  nitric  acid  and  silver  nitrate  added, 
whereby  AgIO4  is  formed,  which  crystallises  out  on  concentration  — 


J2Na2H3I0 

\      2NaI04  +  2AgN03=2NaN03+2AgIO4. 


(2.)  Periodic  acid  is  also  formed  by  the  addition  of  iodine  to 
an  aqueous  solution  of  perchloric  acid  — 

2HC1O4  +  2H2O  +  I2=C12  +  2HIO4,2H2O. 

Properties.  —  The  acid  having-  the  composition  HIO4,2H2O 
is  a  colourless,  crystalline,  deliquescent  substance.  It  melts  at 
133°,  and  at  150°  is  decomposed  into  iodine  pentoxide,  water,  and 
oxygen  — 

2H5IO0=I2O5  +  5H2O  +  O2. 

The  acid  cannot  be  converted  into  HIO4  by  heat,  for  oxygen  is 
evolved  as  soon  as  water  begins  to  be  given  off. 

The  Periodates  constitute  a  numerous  class  of  salts,  many  of  them  being  of 
a  highly  complex  composition.  On  the  assumption  that  iodine  is  monovalent 
in  these  compounds,  their  classification  is  somewhat  difficult,  and  they  must  be 
represented  as  associations  of  molecules  of  salts  of  the  unknown  monobasic 
periodic  acid,  HIO4,  with  metallic  oxide  and  water  in  various  proportions  — 
thus,  the  silver  periodate  in  the  foregoing  equation,  Ag2H3IO6,  would  be 
expressed  by  the  formula,  2AgIO4,Ag2O,2H2O. 

The  classification  of  these  compounds  is  much  simplified  if  we  regard  iodine 
as  here  functioning  as  a  heptavalent  element.  On  this  assumption  the  perio- 
dates  may  be  considered  as  the  salts  of  various  hypothetical  acids,  which  are 
all  derived  from  the  compound  I(HO)7  (itself  hypothetical)  by  the  withdrawal 
of  varying  quantities  of  water.  Thus,  by  the  successive  removal  of  one  mole- 
cule of  water,  the  following  three  acids  would  be  formed— 

I(HO)7  -H20  =  IO(HO)5  .  .  .  H5I06.  (i.) 
IO(HO)5  -H20  =  IO.,(HO)3  .  .  .  H3I05.  (2.) 
I02(HO)3  -H2O=I03(HO)  .  .  .  HIO4.  (3.) 

From  these  three  acids  the  following  salts  may  be  regarded  as  being 
derived  — 


(i.)  Na^H^Oe;  AgoH3IO6  ;  Ag5IO6  ;  Ba5(IO6}2. 
(2.)  Ag3I05;  Pb3(I05)o. 
(3.)  KI04;  AgI04. 


Hypoiodous  Acid  and  Hypoiodites  395 

By  the  abstraction  of  one  molecule  of  water  from  two  molecules  of  these 
acids,  still  more  complex  acids  would  ba  derived,  thus  — 

IO(HO)4 


IO(HO)4 
I02(HO2 


I02(HO)2 

/And  from  these  two  acids  the  following  periodates  may  be  regarded  as  being 
derived  — 

(4.)  Zn4I2On;  Ba4I2On. 
(5.)  Ag4I,09; 


HYPOIODOUS   ACID    AND    HYPOIODITES. 

When  an  aqueous  solution  of  iodine  is  added  to  either  ammonia,  potassium, 
or  sodium  hydroxides,  lime-water  or  baryta-water,  a  colourless  solution  is 
obtained  which  possesses  bleaching  properties.  The  liquid  is  a  dilute  solution 
of  the  hypoiodite  and  iodide  of  the  alkali  used.  Somewhat  stronger  solutions 
may  be  produced  by  adding  small  quantities  of  powdered  iodine  to  the 


A  dilute  solution  of  the  acid  itself  is  obtained  by  shaking  mercuric  oxide 
with  iodine  and  water  (see  Hypochlorous  Acid,  p.  373). 

The  solution  of  the  alkaline  hypoiodite  obtained  by  the  above  reaction  pos- 
sesses well-marked  bleaching  properties.  When  freshly  prepared  it  is  without 
action  upon  starch,  but  is  immediately  decomposed  by  even  so  feeble  an  acid 
as  carbonic  acid,  when  the  blue  starch  compound  is  at  once  formed. 

A  compound  of  iodine  with  lime,  analogous  to  bleaching  powder,  has  been 
obtained  by  shaking  powdered  iodine  with  milk  of  lime.  The  compound  in 
the  presence  of  water  appears  to  behave  in  the  same  way  as  bleaching  powder, 
yielding  a  solution  of  calcium  hypoiodite  and  calcium  iodide  — 

2Ca(OI)I  =  Ca(OI)2+CaI2. 

On  filtering  the  mixture  a  colourless  liquid  is  obtained,  which  gives  no  reaction 
with  starch,  but  which  yields  iodine  when  treated  with  an  acid. 

Neither  the  acid  nor  any  of  its  salts  has  been  isolated,  being  known  only  in 
dilute  solution.  The  compounds  are  all  extremely  unstable,  decomposing  at 
the  ordinary  temperature  in  a  few  hours,  and  in  a  few  minutes  when  the 
solutions  are  boiled  ;  the  salts  passing  into  iodides  and  iodates  — 

3KIO=2KI  +  KIO3, 

while  the  acid  decomposes  first  into  hydriodic  and  iodic  acids,  which   then 
react  upon  each  other  with  elimination  of  free  iodine. 


396  Inorganic  Chemistry 


COMPOUNDS  OF  THE  HALOGENS  WITH  EACH  OTHER. 

Chlorine  unites  both  with  bromine  and  with  iodine,  and  the  two  latter 
elements  combine  with  each  other. 

(i.)  Chlorine  and  Bromine.  —  Bromine  monochloride.  This  substance  is 
obtained  as  a  reddish-  yellow  liquid,  when  chlorine  gas  is  passed  into  bromine. 
The  compound  is  believed  to  have  the  composition  BrCl. 

(2.)  Chlorine  and  Iodine.  —  Iodine  monochloride,  IC1.  When  dry  chlorine 
is  passed  over  iodine,  the  latter  rapidly  melts,  forming  a  dark  reddish-brown 
liquid,  strongly  resembling  bromine  in  appearance.  The  liquid  solidifies  to  a 
mass  of  red  prismatic  crystals,  which  melt  at  25°.  It  is  decomposed  by  water 
intoiodic  and  hydrochloric  acids,  and  iodine  is  liberated  — 

5ICl-t-3H2O=HIO3  +  5HCl  +  2I2. 

Iodine  trichloride,  IC13.  This  compound  is  formed  by  passing  an  excess  of 
chlorine  over  iodine,  or  by  passing  chlorine  through  iodine  monochloride.  It 
is  also  formed  when  hydriodic  acid  is  acted  upon  by  an  excess  of  chlorine  — 


Iodine  trichloride  is  a  yellow  solid  substance,  crystallising  in  long  brilliant 
needle-shaped  crystals,  which  sublime  at  the  ordinary  temperature.  When 
gently  warmed  it  melts,  at  the  same  time  dissociating  into  chlorine  and  the 
monochloride  ;  on  cooling,  reunion  takes  place  with  the  reformation  of  IC13. 

(3.)  Bromine  and  Iodine.  —  Two  compounds  of  these  elements  are  believed 
to  exist,  viz.,  a  crystalline  solid  and  a  deep-coloured  liquid.  Their  composition 
is  probably  expressed  by  the  formulas,  IBr  and  IBrs. 


CHAPTER  II 
THE  ELEMENTS   OF   GROUP    VI.   (FAMILY  B.) 

Oxygen,  O      .         .     16.00      I      Selenium,  Se      .         .     79.2 
Sulphur,  S      .         .     32.07      I      Tellurium,  Te    .         .   127.5 

THE  relation  in  which  oxygen,  the  typical  element,  stands  to  the 
remaining  members  of  the  family  is  very  similar  to  that  between 
fluorine  and  the  other  halogens. 

All  the  elements  of  this  family  unite  with  hydrogen,  forming 
compounds  of  the  type  RH2 — 

OH2,     SH,,     SeH2,     TeH2 ; 

but  the  hydride  of  oxygen  stands  apart  from  the  others  in  many  of 
its  attributes.  Thus  at  ordinary  temperatures  it  is  a  colourless  and 
odourless  liquid,  while  the  remaining  compounds  are  all  foetid- 
smelling  and  poisonous  gases. 

Sulphur,  selenium,  and  tellurium  each  combines  with  oxygen, 
forming  respectively  SO3,  SeO3,  and  TeO3,  while  none  of  these 
elements  in  a  divalent  capacity  forms  a  similar  compound  ;  that  is 
to  say,  no  such  combinations  are  known  as  OS3,  or  OSe3,  although 
amongst  themselves  they  unite,  forming  SeS2  and  TeS2. 

Sulphur,  selenium,  and  tellurium  also  unite  with  oxygen,  forming 
dioxides,  SO2,  SeO9,  and  TeO2,  in  which  these  elements  are  pos- 
sibly tetravalent,  in  which  case  the  constitution  of  the  compounds 
will  be  represented  thus,  O  =  S  =  O  ;  O  =  Se  =  O. 

We  may,  however,  consider  them  as  functioning  in  a  divalent 

/O  /O 

capacitv,  and  regard  the  oxides  as  constituted  thus,  S<(     j  ;  Se<(     I  < 

\O          NO 
in  which  case  we  may  look  upon  ozone  as  being  the  corresponding 

/° 
oxygen  compound,  OO2,  O<^    |  . 

357 


398  Inorganic  Chemistry 

All  the  elements  of  this  family  combine  with  chlorine,  producing 
compounds  having  the  following  composition  : — 

Oxygen.  Sulphur.  Selenium.  Tellurium. 

02C1 

—  SoCLj  SeaClg  — 

OC12  SC12  TeCl2 

SQ4  SeCl4  TeCl4 

Oxygen  again  differs  from  the  other  members  by  alone  forming 
a  compound  of  the  type,  R2C1.  This  element  also  shows  no  ten- 
dency to  function  with  a  higher  atomicity  than  that  of  a  divalent ; 
while  the  others  unite  with  four  atoms  of  the  halogen,  thereby 
exhibiting  their  tetravalent  nature. 

The  members  of  this  family  pass  by  a  regular  gradation  from 
the  strongly  electro-negative,  gaseous,  non-metal  oxygen  to  the 
feebly  negative  and  slightly  basic  element  tellurium,  which  possesses 
many  of  the  properties  of  a  true  metal.  Selenium  and  tellurium 
are  both  elements  which  lie  very  close  to  that  ill-defined  boundary 
between  the  metals  and  non-metals,  and  are  on  this  account  some- 
times termed  metalloids.  In  tellurous  oxide,  TeO2,  we  have  a 
compound  which  is  both  an  acid-forming  and  a  salt-forming  oxide, 
its  acidic  and  basic  properties  being  nearly  equally  balanced.  Thus, 
it  unites  with  water,  forming  tellurous  acid,  H2TeO3,  corresponding 
to  sulphurous  acid,  H2SO3  j  while  tellurium  replaces  hydrogen  in 
sulphuric  acid,  forming  tellurium  sulphate,  Te(SO4)2. 

Of  the  four  elements  of  this  family,  oxygen  is  by  far  the  most 
abundant,  both  in  combination  and  in  the  free  state  ;  sulphur  is 
more  plentiful  than  the  other  two,  and  tellurium  occurs  in  the 
smallest  quantity. 

The  element  oxygen  has  already  been  treated  in  Part  II. 


SULPHUR. 

Symbol,  S.     Atomic  weight  =  32. 07.     Molecular  weight =64. 14. 

Occurrence.— In  the  free  state  this  element  occurs  chiefly  in 
volcanic  districts.  In  Italy  and  Sicily  large  quantities  of  native 
sulphur  are  found,  which  have  long  been  the  most  important 
European  sources  of  this  substance.  Large  deposits  are  to  be  met 
with  in  Transylvania  and  in  Iceland,  and  it  also  occurs  in  beds, 


Sulphur  399 

often  of  great  thickness,  in  parts  of  China,  India,  California,  and 
the  Yellowstone  district  of  the  Rocky  Mountains.  These  natural 
deposits  are  sometimes  found  stratified  with  beds  of  clay  or  rock, 
but  they  often  occur  as  what  are  known  as  •'  living  "  beds,  in  which 
the  sulphur  is  continuously  being  formed  as  the  result  of  chemical 
decompositions  which  are  at  present  at  work.  Such  a  "  living  " 
sulphur  bed  is  known  as  a  solfatara,  and,  as  in  the  case  of  the 
Iceland  deposits,  they  are  usually  found  associated  with  geysers, 
fumaroles,  and  other  signs  of  volcanic  action. 

In  combination  with  hydrogen,  sulphur  occurs  as  sulphuretted 
hydrogen.  Enormous  quantities  of  sulphur  are  found  combined 
with  various  metals,  constituting  the  important  class  of  substances 
known  as  sulphides  ;  as,  for  example,  galena,  or  lead  sulphide, 
PbS  ;  zinc  blende,  or  zinc  sulphide,  ZnS  ;  Pyrites,  or  iron  sulphide, 
FeSa  ;  copper  pyrites,  or  copper  iron  sulphide,  Cu2Fe2S4  ;  stibnite^ 
or  antimony  sulphide,  SboSs  ;  cinnabar,  or  mercury  sulphide,  HgS. 

In  combination  with  metals  and  oxygen,  sulphur  occurs  in 
sulphates,  such  as  gypsum,  CaSO4,2H2O  ;  heavy  spar,  BaSO4  ; 
kieserite,  MgSO4,H2O. 

Modes  Of  Formation.—  (  i.)  Sulphur  is  formed  when  sulphu- 
retted hydrogen  is  brought  in  contact  with  sulphur  dioxide  ;  the 
two  gases  mutually  decompose  one  another  with  the  formation  of 
water  and  the  precipitation  of  sulphur  — 

2H2S  +  SO2=2H2O  +  3S. 

(2.)  It  is  also  produced  when  sulphuretted  hydrogen  is  burnt 
with  an  insufficient  supply  of  air  — 


This  reaction  probably  takes  place  in  two  stages,  a  portion 
of  the.  sulphuretted  hydrogen  burning  to  sulphur  dioxide,  and  this 
then  reacting  upon  a  further  quantity  of  sulphuretted  hydrogen, 
thus— 

(a)  H2S  +  3O 

(£)  2H2S  +  SO2=2H2 


It  is  supposed  that  some  of  the  free  sulphur  found  in  volcanic 
regions  has  been  produced  by  this  action  of  these  two  gases  upon 
one  another. 

Extraction  of  Sulphur  from  Native  Sulphur.—  Natural 
sulphur  is  always  more  or  less  mixed  with  earthy  or  mineral 


400  Inorganic  Chemistry 

matters,  from  which  it  is  necessary  to  free  it.  This  is  usually 
effected  by  melting  the  sulphur  and  allowing  it  to  flow  away  from 
the  accompanying  impurities.  The  crude  sulphur  rock  is  stacked 
in  brick  kilns  having  a  sloping  floor,  and  the  mass  ignited  by 
introducing  through  openings  in  the  heap  burning  faggots  of 
brushwood.  The  heat  produced  by  the  combustion  of  a  part  of  the 
sulphur  causes  the  remainder  to  melt  and  collect  upon  the  sloping 
floor  of  the  kiln,  from  which  it  can  be  drawn  off  into  rough  moulds. 
The  loss  of  sulphur  by  this  method  is  very  considerable,  usually 
not  more  than  two-thirds  of  the  total  amount  contained  in  the  rock 
being  obtained. 

(3.)  Sulphur  may  be  obtained  by  heating  certain  metallic  sul- 
phides ;  thus  when  iron  pyrites  is  heated  it  yields  one-third  of  its 
sulphur  — 

2  =  Fe3S4+S2. 


If  the  pyrites  be  roasted  in  kilns,  the  whole  of  the  sulphur  is 
obtained,  partly  as  free  sulphur  and  partly  as  sulphur  dioxide  — 


This  method  was  at  one  time  rather  extensively  employed  for 
the  preparation  of  sulphur  on  a  manufacturing  scale,  but  has  now 
practically  gone  out  of  use,  the  pyrites  being  usually  roasted  with 
excess  of  air,  whereby  the  whole  of  the  sulphur  is  converted  into 
sulphur  dioxide  for  use  in  the  manufacture  of  sulphuric  acid. 

By  a  similar  process,  sulphur  is  obtained  as  a  bye-product  during 
the  roasting  of  copper  pyrites  in  the  first  stage  of  the  operation  of 
copper-smelting. 

(4.)  Large  quantities  of  sulphur  are  now  extracted  from  the  vat- 
waste  or  alkali-waste,  obtained  in  the  manufacture  of  sodium 
carbonate  by  the  Leblanc  process.  This  material  consists  largely 
of  an  insoluble  oxy-sulphide  of  lime,  a  compound  containing  calcium 
sulphide  (CaS)  and  calcium  oxide  (CaO)  in  varying  proportions. 
Moneys  process^  which,  however,  has  now  been  entirely  superseded 
by  Chance's  process  (p.  411),  is  the  following  :  A  current  of  air  is 
blown  through  the  compound,  whereby  the  calcium  sulphide  it 
contains  is  ultimately  converted  into  a  mixture  of  calcium  hydro- 
sulphide  (CaH2S2),  thiosulphate  (CaS2O3),  and  polysulphide  (CaS5), 
according  to  the  following  equations  — 

O  =  CaHS  +  CaHO. 


Sulphur  401 

This  reaction  goes  forward  in  several  stages,  in  the  course  of 
which  a  quantity  of  sulphur  is  set  free  ;  this  is  then  acted  upon  by 
the  calcium  hydroxide,  with  the  formation  of  calcium  polysulphide 
and  calcium  thiosulphate,  thus  — 

(2.)  3CaH2O2  +  12S  =  2CaS5  +  CaS2O3  +  3H2O. 

The  materials  are  alternately  oxidised  and  lixiviated  several 
times,  and  the  liquor  is  then  treated  with  excess  of  hydrochloric 
acid  at  a  temperature  of  about  60°,  which  decomposes  the  various 
sulphur  compounds  according  to  the  following  equations  — 

(a.) 
(J.) 

(c.) 

The  best  results  are  obtained  when  the  sulphur  compounds  are 
present  in  such  proportions  that  the  SO2  evolved  by  reaction  c  is 
sufficient  to  decompose  the  whole  of  the  SH2  produced  by  the 
other  two  reactions,  so  that  neither  gas  escapes  — 


(5.)  Sulphur  is  also  obtained  from  the  spent  oxide  of  iron  which 
has  been  used  in  the  "  purifiers"  employed  upon  gas-works.  Coal 
gas  contains  considerable  quantities  of  sulphuretted  hydrogen, 
which  are  removed  from  the  gas  by  passing  it  through  hydrated 
ferric  oxide  (Fe2H6O6),  which  absorbs  the  whole  of  the  sulphuretted 
hydrogen,  thus  — 

Fe2H6O6  +  3H2S  =  2FeS  +  S  +  6H2O. 

When  the  compound  has  lost  its  power  to  absorb  sulphuretted 
hydrogen,  the  material  is  thrown  out  of  the  purifiers  and  exposed 
to  air  and  moisture,  when  the  iron  becomes  reconverted  into  the 
hydrated  oxide,  and  the  sulphur  is  set  free  — 

2FeS  +  3O  +  3H2O  =  Fe2H6O6  +  2S. 

This  revivified  material  is  then  employed  for  the  purification  of 
a  further  quantity  of  gas.  It  is  found  that  after  a  certain  number 
of  revivifying  operations  the  substance  begins  to  lose  its  power  of 
absorbing  any  additional  sulphuretted  hydrogen,  and  as  it  then 

2  C 


402 


Inorganic  Chemistry 


contains  nearly  half  its  weight  of  sulphur,  it  becomes  a  valuable 
source  of  this  element.  The  sulphur  is  obtained  from  it  by  distil- 
lation, or  the  material  may  be  roasted  in  special  kilns,  whereby  the 
sulphur  is  converted  into  sulphur  dioxide,  and  employed  for  the 
manufacture  of  sulphuric  acid. 

Purification. — The  crude  sulphur  obtained  by  the  foregoing 
methods  is  purified  by  distillation,  the  process  being  carried  out  in 
the  arrangement  shown  in  Fig.  109.  The  sulphur  is  first  melted 
in  an  iron  pot  d,  and  the  liquid  substance  drawn  off  at  will  by 


FIG.  109. 

means  of  the  pipe  F  into  the  retort  B.  The  sulphur  is  there 
boiled  by  means  of  the  fire,  and  the  vapour  allowed  to  issue  into  the 
large  brickwork  chamber  G.  As  the  vapour  enters  the  chamber,  it 
condenses  upon  the  walls  and  floor  in  the  form  of  a  light,  powdery 
deposit,  consisting  of  minute  crystals,  and  constituting  the  flowers 
of  sulphur  of  commerce.  As  the  process  continues,  and  the  brick- 
work becomes  hot,  this  soft  powder  melts  and  collects  upon  the 
floor  as  an  amber- coloured  liquid,  which  is  run  out  from  time  to 


Sulphur  403 

time  from  the  opening  at  H,  and  cast  either  into  large  blocks  or 
into  cylindrical  rods,  by  means  of  wooden  moulds.  In  the  latter 
form  it  is  known  as  roll  sulphur. 

When  the  sulphur  vapour  first  enters  the  chamber  and  mixes 
with  the  air,  the  mixture  frequently  ignites  with  a  feeble  explosion  ; 
the  chamber,  therefore,  is  furnished  with  a  valve,  S,  at  the  top, 
whereby  the  pressure  developed  at  the  moment  of  combustion  may 
be  relieved. 

Properties. — Sulphur,  as  ordinarily  seen,  is  a  pale-yellow  brittle 
crystalline  solid.  It  is  insoluble  in  water,  but  readily  dissolves  in 
carbon  disulphide,  and  to  a  greater  or  less  degree  in  turpentine, 
benzene,  chloroform,  sulphur  chloride,  and  many  other  solvents. 
It  is  a  non-conductor  of  electricity,  and  an  extremely  bad  con- 
ductor of  heat.  A  piece  of  sulphur  on  being  very  gently  warmed, 


FIG.  no. 

even  by  being  grasped  in  the  hand,  may  be  heard  to  crack  by  the 
unequal  heating,  and  will  ultimately  fall  to  pieces.  At  a  tem- 
perature of  114.5°  sulphur  melts  to  a  clear  amber-coloured  and 
moderately  mobile  liquid ;  on  raising  the  temperature  of  this 
liquid  its  colour  rapidly  darkens,  and  at  the  same  time  it  loses  its 
mobility,  until  at  a  temperature  of  about  230°  the  mass  appears 
almost  black,  and  is  so  viscous  that  it  can  no  longer  be  poured 
from  the  vessel.  As  the  temperature  is  still  further  raised,  the 
substance,  while  retaining  its  dark  colour,  again  becomes  liquid, 
although  it  does  not  regain  its  original  limpidity.  At  448°  the 
liquid  boils,  and  is  converted  into  a  pale  yellowish-brown  coloured 
vapour.  On  allowing  the  boiling  sulphur  to  cool,  it  passes  through 
the  same  changes  in  reverse  order  until  it  solidifies. 

When  the  vapour  of  sulphur  is  heated  to  1000°,  it  is  converted 


404 


Inorganic  Chemistry 


into  a  true  gas,  and  has  a  density  of  32,  one  litre  of  the  gas  weigh- 
ing 32  criths. 

Sulphur  is  known  to  exist  in  four  allotropic  modifications,  two  of 
which  are  crystalline  and  two  amorphous. 

(a)  "Rhombic"  Sulphur.— Of  the  two  crystalline  varieties  this  is 
the  more  stable.  Sulphur,  therefore,  that  occurs  native  is  found 
crystallised  in  this  form,  namely,  orthorhombic  pyramids.  It  may 
be  obtained  by  allowing  a  solution  of  sulphur  in  carbon  disulphide 
to  slowly  evaporate.  Fig.  no  represents  two  large  crystals  of 
sulphur  obtained  in  this  way. 

Orthorhombic  crystals  of  sulphur  can  also  be  obtained  under 
certain  conditions  when  melted  sulphur  is  allowed  to  crystallise. 


FIG.  in. 

Sulphur  in  the  liquid  condition  exhibits  the  phenomenon  of  sus- 
pended solidification  to  a  very  high  degree,  and  if  the  liquid  be 
carefully  cooled  out  of  contact  with  air,  the  temperature  may  fall 
to  90°  before  solidification  takes  place.  If  into  the  liquid  in  this 
state  a  crystal  of  the  rhombic  variety  be  dropped,  the  sulphur 
begins  to  solidify  in  crystals  of  that  form.  If  the  superfused  sulphur 
be  contained  in  a  hermetically  closed  flask,  the  liquid  frequently 
deposits  orthorhombic  crystals,  and  by  allowing  the  mass  to 
partially  solidify,  and  quickly  inverting  the  flask,  the  crystals  may 
be  seen  upon  the  bottom  of  the  vessel. 

The  specific  gravity  of  this  form  of  sulphur  is  2.05. 

(/3)  "Prismatic"  Sulphur. — When   melted  sulphur   is   allowed 


Sulphur  405 

to  cool  under  ordinary  conditions,  such  as  in  a  crucible  or 
beaker,  it  crystallises  in  the  form  of  prismatic  needles,  belong- 
ing to  the  monoclinic  or  monosymmetric  system.  By  allowing 
the  mass  to  partially  solidify,  and  pouring  off  the  still  liquid  por- 
tion, these  crystals  will  be  seen  lining  the  inside  of  the  beaker 
as  long  translucent  prisms.  Fig.  ill  shows  such  a  mass  of 
crystals.  Prismatic  crystals  of  sulphur  are  also  obtained  when 
this  element  is  crystallised  from  a  hot  solution  in  oil  of  tur- 
pentine. 

The  specific  gravity  of  this  form  of  sulphur  is  less  than  that  of 
the  orthorhombic  variety,  being  1.98. 

At  ordinary  temperatures  this  modification  is  unstable,  and  in 
the  course  of  a  day  or  two  the  crystals  lose  their  translucent 
appearance,  owing  to  their  becoming  broken  down  into  a  number 
of  smaller  crystals  of  the  rhombic  variety,  and  present  the  opaque 
yellow  appearance  of  ordinary  roll  sulphur.  This  change  from 
the  prismatic  to  the  "  rhombic  "  variety,  which  takes  place  more 
quickly  when  the  crystals  are  scratched  or  subjected  to  vibration, 
is  attended  with  evolution  of  heat.  When  monoclinic  sulphur  is 
thrown  into  carbon  disulphide,  its  transformation  into  the  stable 
modification  takes  place  rapidly,  and  in  this  way,  by  means  of  a 
thermopile,  the  heat  evolved  by  the  change  may  be  rendered 
evident.  As  carbon  disulphide,  however,  at  once  exerts  its  solvent 
action  upon  the  "rhombic"  sulphur  the  moment  it  is  formed,  the 
reduction  of  temperature  resulting  from  this  cause  would  com- 
pletely overbalance  and  mask  the  more  feeble  heat  effect  produced 
by  the  passage  of  the  sulphur  from  the  unstable  to  the  stable  form. 
In  order,  therefore,  to  render  evident  the  heat  resulting  from  the 
change  of  crystalline  form,  the  carbon  disulphide  must  be  pre- 
viously allowed  to  dissolve  as  much  sulphur  as  it  can  take  up.  If 
a  small  quantity  of  carbon  disulphide,  so  saturated  with  sulphur, 
be  placed  in  a  corked  flask,  and  stood  upon  the  face  of  a  thermo- 
electric pile*  in  connection  with  a  galvanometer,  and  a  quantity 
of  prismatic  crystals  of  sulphur  be  quickly  thrown  into  the  liquid, 
a  sensible  deflection  of  the  galvanometer  needle  will  be  seen  in  the 
direction  caused  by  heat. 

Although  under  ordinary  conditions  monoclinic  sulphur  is  un- 
stable and  passes  into  the  "  rhombic  "  form,  at  temperatures  between 

*  The  thermo-electric  pile  is  a  delicate  physical  instrument  employed  for 
registering  slight  changes  of  temperature;  for  descriptions  of  the  apparatus 
the  student  must  consult  text-books  on  physics. 


406 


Inorganic  Chemistry 


100°  and  114°  it  appears  to  be  the  more  stable  variety,  for  at  this 
temperature  "rhombic  "  sulphur  passes  into  the  monoclmic  variety. 
(y)  Plastic  Sulphur.— When  sulphur  which  has  been  heated 
-until  it  reaches  the  viscous  condition  is  suddenly  plunged  into 
water,  or  when  boiling  sulphur  is  poured  in  a  thin  stream  into 
water,  the  substance  solidifies  to  a  tough  elastic  material  some- 
what resembling  indiarubber.  The  sulphur  in  this  form  is  known 
as  plastic  sulphur.  This  variety  is  best  obtained  by  distilling  a 
quantity  of  ordinary  sulphur  from  a  glass  retort  (Fig.  112),  and 
allowing  the  distilled  liquid  to  flow  in  a  fine  stream  into  cold 
water  placed  for  its  reception.  As  the  liquid  sulphur  falls  into 
the  water,  it  congeals  to  the  plastic  condition  as  a  continuous 

thread,  which  winds  itself  in 
a  regular  manner  into  beauti- 
ful coils  of  a  delicate  trans- 
lucent amber  colour.  The 
specific  gravity  of  plastic  sul- 
phur is  1.95,  and  it  is  not 
soluble  in  carbon  disulphide. 
At  ordinary  temperatures  this 
allotrope  of  sulphur  is  gra- 
dually transformed  into  the 
stable  "rhombic"  variety  ;  in 
the  course  of  a  few  days  it 
FIG.  112.  loses  its  transparency  and 

elasticity,  and  becomes  con- 
verted into  the  ordinary  lemon-yellow  brittle  condition  of  common 
sulphur.  This  change  takes  place  more  quickly  if  the  plastic 
material  be  stretched  and  worked  between  the  fingers,  and  still 
more  readily  by  heating  it  for  a  few  moments  to  100°,  and  allowing 
it  again  to  cool. 

(8)  White  Amorphotis  Sulphur. — When  sulphur  is  heated,  and 
the  vapour  condensed  upon  a  cool  surface,  as  in  the  formation  of 
ordinary  flowers  of  sulphur,  although  the  greater  portion  of  the 
sulphur  is  sublimed  in  the  orthorhombic  form,  the  sublimate  con- 
tains a  small  amount  of  sulphur  in  the  form  of  an  amorphous 
powder,  which  is  almost  milk-white  in  colour. 

This  modification  is  best  obtained  by  treating  flowers  of  sulphur, 
which  usually  contains  as  much  as  5  or  6  per  cent,  of  amorphous 
sulphur,  with  carbon  disulphide,  whereby  the  orthorhombic  variety' 
is  dissolved,  and  the  white  amorphous  substance,  which  is  insoluble; 


Sulphur  407 

in  that  liquid,  is  left  behind.  By  filtering  the  liquid  and  washing 
the  residue  with  carbon  disulphide  until  the  whole  of  the  soluble 
sulphur  is  removed,  the  amorphous  powder  may  be  obtained  in  a 
state  of  purity. 

This  amorphous  substance  is  also  produced  in  small  quantity,  by 
the  action  of  light  upon  a  solution  of  sulphur  in  carbon  disulphide. 
Thus,  if  a  perfectly  clear  solution  of  sulphur  in  this  liquid  be  placed 
for  even  a  few  minutes  in  the  path  of  a  beam  of  electric  light,  the 
solution  will  be  seen  to  become  rapidly  turbid,  owing  to  the  forma- 
tion of  this  insoluble  modification. 

This  milk-white  amorphous  modification  is  stable  at  the  ordinary 
temperature,  and  therefore  does  not  pass  spontaneously  into  the 
rhombic  variety.  When  heated  to  a  temperature  of  100°,  it  quickly 
becomes  yellow  in  colour,  and  is  then  readily  soluble  in  carbon 
disulphide,  having  been  transformed  at  that  temperature  into  the 
ordinary  stable  form. 

Milk  Of  Sulphur.  —  This  substance  is  a  medicinal  preparation, 
obtained  by  precipitating  sulphur  from  a  polysulphide  of  lime  by 
means  of  hydrochloric  acid.  Flowers  of  sulphur  and  milk  of  lime 
are  boiled  together  for  some  time,  and  after  settling  the  clear 
reddish  liquid  containing  the  calcium  polysulphides  is  decanted  off, 
and  hydrochloric  acid  added  to  it  ;  calcium  chloride  is  formed, 
and  sulphur  in  a  fine  state  of  subdivision  is  precipitated  — 


The  product  so  obtained  is  pale  yellow  in  colour,  and  consists  of 
ordinary  sulphur  often  contaminated  with  considerable  quantities 
of  calcium  sulphate,  derived  from  sulphuric  acid  present  in  the 
hydrochloric  acid  employed  in  the  precipitation. 

When  sulphur  in  any  of  its  modifications  is  heated  in  the  air,  it 
takes  fire  and  burns  with  a  pale  blue  flame,  giving  rise  to  sulphur 
dioxide  ;  when  burnt  in  oxygen  a  small  quantity  of  sulphur  tri- 
oxide  is  at  the  same  time  produced. 

Finely  divided  sulphur,  when  exposed  to  air  and  moisture,  under- 
goes slow  oxidation  even  at  ordinary  temperatures,  with  the  forma- 
tion of  sulphuric  acid.  Thus,  if  flowers  of  sulphur  be  moistened 
with  water  and  freely  exposed  to  the  air,  in  a  short  time  the  water 
will  be  distinctly  acid.  On  this  account  sulphur  that  is  used  for 
pyrotechnic  purposes  is  thoroughly  washed  and  dried,  and  pre- 
served in  warm  dry  places. 


408  Inorganic  Chemistry 

Sulphur  combines  directly  with  many  metals  under  the  influence 
of  heat,  forming  sulphides,  the  union  in  many  cases  being  accom- 
panied by  vivid  combustion.  Thus,  a  strip  of  copper,  when  intro- 
duced into  sulphur  vapour,  burns  brilliantly  with  the  formation  of 
copper  sulphide,  and  a  red-hot  bar  of  iron,  when  pressed  against 
a  roll  of  sulphur,  burns  in  the  vapour  which  is  generated,  and  the 
molten  sulphide  falls  in  scintillating  masses  through  the  air — 

FeS. 


Heated  with  sodium  or  potassium,  the  alkaline  sulphides  are 
formed  with  deflagration — 

K2+S  =  K2S. 

COMPOUNDS  OF  SULPHUR  WITH  HYDROGEN. 

Two  compounds  of  these  elements  are  known,  namely — 
Hydrogen  sulphide  or  sulphuretted  hydrogen         .     H2S. 
Hydrogen  persulphide H2S2. 

HYDROGEN  SULPHIDE. 

Formula,  H2S.     Molecular  weight  =  34.08.     Density  =  17.04.  ' 

Occurrence. — This  gas  is  evolved  in  volcanic  regions,  and  is 
met  with  in  solution  in  sulphur  mineral  waters. 

Modes  of  Formation.— ( i.)  Sulphuretted  hydrogen  may  be 
formed  by  the  direct  union  of  its  elements,  by  passing  a  mixture  of 
hydrogen  and  the  vapour  of  sulphur  through  a  strongly  heated 
tube.  In  small  quantity  it  is  produced  when  hydrogen  is  passed 
into  boiling  sulphur,  or  over  certain  heated  metallic  sulphides. 

(2.)  Sulphuretted  hydrogen  is  most  readily  obtained  by  the 
action  of  either  hydrochloric  or  sulphuric  acid  upon  ferrous  sul- 
phide, thus — 

FeS  +  2HC1  =  FeCl2+  H2S. 
FeS  +  H2SO4= FeSO4  +  H2S. 

The  ferrous  sulphide  in  broken  fragments  is  placed  in  a  two- 
necked  bottle,  similar  to  the  apparatus  (Fig.  29)  employed  for 
the  preparation  of  hydrogen,  and  the  dilute  acid  poured  upon  it, 
The  gas  is  rapidly  evolved  without  the  application  of  heat.  The 
gas  obtained  by  this  method  always  contains  free  hydrogen, 
owing  to  the  presence  of  uncombined  iron  in  the  ferrous  sulphide. 


Hydrogen  Sulphide  409 

(3.)  Pure  hydrogen  sulphide  may  be  obtained  by  heating  anti- 
mony trisulphide  (grey  antimony  ore)  with  strong  hydrochloric  acid — 


(4.)  Also  by  the  action  of  water  upon  aluminium  sulphide 
(p.  623). 

(5.)  Sulphuretted  hydrogen  is  produced  during  the  putrefaction 
of  organic  substances  containing  sulphur,  the  offensive  smell  of  a 
decomposing  egg  being  due  to  the  presence  of  this  gas.  It  is  also 
produced  during  the  destructive  distillation  of  coal,  by  the  direct 
union  of  hydrogen  with  the  sulphur  contained  in  the  pyrites,  hence 
coal  gas  always  contains  sulphuretted  hydrogen  amongst  its  im- 
purities. 

Properties.  —  Sulphuretted  hydrogen  is  a  colourless  gas,  having  a 
somewhat  sickly  sweetish  taste  and  an  extremely  offensive  odoun 
It  acts  as  a  powerful  poison  when  inhaled  in  the  pure  state,  and 
even  when  very  largely  diluted  with  air  it  gives  rise  to  dizziness 
and  headache.  Its  poisonous  effects  are  more  marked  upon  some 
animals  than  others  :  thus,  a  bird  was  found  to  die  in  an  atmosphere 
containing  only  ^^j  of  sulphuretted  hydrogen,  while  it  required  an 
amount  equal  to  gou  to  poison  a  hare  ;  and  again,  cold-blooded 
animals  are  in  no  way  affected  by  inhaling  these  proportions  of 
the  gas.  Sulphuretted  hydrogen  is  moderately  soluble  in  water  ; 
at  ordinary  temperatures  water  dissolves  about  three  times  its  own 
volume  of  the  gas.  In  collecting  it  over  water,  therefore,  consider- 
able loss  results  unless  the  water  be  warm.  The  coefficient  of 
absorption  by  water  at  o°  is  4.3706. 

The  aqueous  solution  gives  an  acid  reaction  with  litmus,  and 
possesses  the  taste  and  smell  of  the  gas.  It  quickly  decomposes 
on  exposure  to  air,  the  hydrogen  of  the  sulphuretted  hydrogen 
combines  with  oxygen,  and  the  liquid  becomes  turbid  by  the  preci- 
pitation of  sulphur.  Hydrogen  sulphide  is  an  inflammable  gas, 
burning  with  a  bluish  flame,  and  producing  sulphur  dioxide  and 
water  — 


If  mixed  with  oxygen  in  the  proportion  demanded  by  this  equa- 
tion, viz.,  two  volumes  of  sulphuretted  hydrogen  and  three  volumes 
of  oxygen,  and  ignited,  the  mixture  explodes  with  violence.  When 
the  gas  is  burned  with  an  insufficient  supply  of  air  or  oxygen  for 
its  complete  combustion,  the  sulphur  is  deposited. 

Sulphuretted  hydrogen  is  decomposed  by  the  halogens,  with  the 


4IO  Inorganic  Chemistry 

deposition  of  sulphur,  and  the  formation  of  the  hydrogen  compound 
of  the  halogen  element  thus  — 


Fluorine,  chlorine,  and  bromine  are  capable  of  bringing  about 
this  decomposition  at  ordinary  temperatures  ;  in  the  case  of  iodine, 
the  reaction  is  attended  with  absorption  of  heat,  which  may  be 
supplied  by  passing  the  mixture  of  iodine  vapour  and  sulphuretted 
hydrogen  through  a  hot  tube,  or  by  causing  the  action  to  take 
place  in  the  presence  of  water.  In  the  latter  case  the  heat  of  solu- 
tion of  the  hydriodic  acid  determines  the  reaction. 

When  passed  into  sulphuric  acid,  reduction  of  the  acid  takes 
place  with  the  precipitation  of  sulphur  — 


Sulphuretted  hydrogen,  therefore,  cannot  be  dried  by  means  of 
sulphuric  acid. 

The  gas  acts  upon  many  metals  with  the  formation  of  sulphides  ; 
thus,  when  potassium  is  heated  in  a  stream  of  hydrogen  sulphide, 
it  readily  burns  and  produces  potassium  hydrosulphide  — 


Such  metals  as  tin,  lead,  silver,  &c.,  are  rapidly  tarnished  in 
contact  with  this  gas.  On  this  account  articles  of  silver,  when 
exposed  to  the  air  of  towns,  quickly  become  covered  with  a  film 
of  sulphide,  which  first  appears  yellowish-brown,  and  gradually 
becomes  black.  The  discoloration  of  a  silver  spoon,  when  intro- 
duced into  an  egg  which  is  partially  decomposed,  is  due  to  the 
same  cause. 

Sulphuretted  hydrogen  also  acts  upon  metallic  salts,  combining 
with  the  metal  to  form  a  sulphide.  The  "  white-lead  "  employed 
in  ordinary  paint  is  gradually  blackened  on  prolonged  exposure 
to  the  air  by  the  formation  of  lead  sulphide. 

Hydrogen  sulphide  is  rapidly  absorbed  by  lime,  with  the  forma- 
tion of  calcium  hydrosulphide  — 

CaH2O,  +  2H2S  =  CaH2S2+  2H2O. 

It  is  also  absorbed  by  calcium  sulphide,  yielding  the  same 
compound.  This  reaction  is  employed  in  the  method  known  as 


Hydrogen  Sulphide  411 

Chance  s  process,  for  utilising  the  sulphur  of  the  vat-waste  of  the 
alkali  manufacture.  This  consists  in  passing  lime-kiln  gases 
through  a  series  of  vessels  containing  the  waste  mixed  with  water. 
In  the  first  vessels  the  carbon  dioxide  is  absorbed,  and  sulphuretted 
hydrogen  evolved.  This,  passing  into  the  later  vessels,  is  absorbed 
by  the  vat-waste,  forming  calcium  hydrosulphide,  which  in  its 
turn  is  decomposed  by  carbon  dioxide,  with  the  evolution  of  twice 
the  volume  of  sulphuretted  hydrogen  for  a  given  volume  of  carbon 
dioxide,  as  in  the  first  reaction  — 


(i) 

(2)  CaS  +  H2S=CaH2S 

(3) 


The  sulphuretted  hydrogen,  mixed  with  atmospheric  nitrogen 
and  a  small  quantity  of  carbon  dioxide,  is  sufficiently  rich  to  burn, 
yielding  sulphur  dioxide,  which  can  then  be  employed  for  the 
manufacture  of  sulphuric  acid. 

Sulphuretted  hydrogen  is  also  decomposed  by  ferric  hydroxide> 
with  the  formation  of  ferrous  sulphide  and  water,  and  the  deposi- 
tion of  sulphur,  as  described  on  page  401.  This  action  takes 
place  with  the  evolution  of  considerable  heat,  the  temperature 
rising  high  enough  to  ignite  a  mixture  of  sulphuretted  hydrogen 
and  oxygen.* 

Sulphuretted  hydrogen  is  a  valuable  laboratory  reagent,  on 
account  of  the  general  behaviour  of  certain  classes  of  sulphides. 
Thus,  the  sulphides  of  certain  metals,  being  insoluble  in  dilute 
acids,  are  precipitated  from  acid  solutions  ;  for  example  — 


CdCl2  +H2S  =  CdS  +  2HCl. 

Others  are  soluble  in  acids,  but  insoluble  in  alkaline  liquids,  and 
are  therefore  precipitated  by  sulphuretted  hydrogen  in  the  presence 
of  ammonia,  or  by  the  addition  of  ammonium  sulphide,  thus  — 

ZnSO4  +  (NH4)2S  =  ZnS  +  (NH4)2SO4. 

A  third  group  of  metals  yield  sulphides  that  are  soluble  in  water, 

and  therefore  are  not  separated  either  in  acid  or  alkaline  solutions. 

Many  of  the  metallic  sulphides  are  also  possessed  of  charac- 

"  Chemical  Lecture  Experiments,"  new  ed.,  Nos.  577,  578. 


412  Inorganic  Chemistry 

teristic  colours,  which  readily  serve  for  their  identification.  Thus 
arsenious  sulphide  is  pale  yellow,  and  cadmium  sulphide  golden 
yellow.  Antimonious  sulphide  has  a  bright  red  colour,  while  zinc 
sulphide  is  white. 

This  behaviour  of  metals  towards  sulphuretted  hydrogen  is 
the  basis  upon  which  certain  methods  of  qualitative  analysis  are 
founded. 


HYDROGEN  PERSULPHIDE. 

Formula,  H2S2. 

Modes  Of  Formation.  —  (i.)  This  substance,  which  stands  in  the 
same  relation  to  hydrogen  sulphide  as  hydrogen  peroxide  does  to 
water,  may  be  obtained  by  slowly  pouring  a  solution  of  calcium 
or  sodium  disulphide  into  diluted  hydrochloric  acid  cooled  by  a 
freezing  mixture,  the  liquids  being  rapidly  stirred  during  the  pro- 
cess of  mixing,  and  the  acid  being  kept  in  considerable  excess  — 


The  hydrogen  persulphide  separates  out  as  a  heavy,  pale-yellow, 
oily  compound,  which  settles  to  the  bottom  of  the  liquid. 

Properties.  —  Hydrogen  persulphide  or  hydrogen  disulphide  is 
an  oily  liquid  having  a  specific  gravity  of  1.376.  It  has  a  pungent 
smell,  accompanied  by  the  odour  of  sulphuretted  hydrogen,  due 
probably  to  the  partial  decomposition  of  the  compound,  and  its 
vapour  is  irritating  to  the  eyes.  It  is  an  unstable  substance, 
decomposing  at  ordinary  temperatures  into  sulphur  and  sulphu- 
retted hydrogen  :  when  heated  this  decomposition  takes  place 
rapidly.  It  is  immediately  decomposed  by  alkalis,  but  is  more 
stable  in  the  presence  of  dilute  hydrochloric  acid.  It  is  in- 
soluble in  water,  but  dissolves  readily  in  carbon  disulphide, 
the  solution  in  this  liquid  being  more  stable  than  the  liquid 
substance  itself. 

Hydrogen  persulphide  burns  with  a  blue  flame,  yielding  sulphur 
dioxide  and  water.  It  possesses  feeble  bleaching  properties,  and, 
like  its  oxygen  analogue,  it  is  decomposed  by  certain  metallic 
oxides,  with  the  evolution  of  sulphuretted  hydrogen. 


Hydrogen  Per  sulphide  413 

Hydrogen  persulphide  readily  dissolves  sulphur,  and  owing  to 
the  fact  that  sulphur  is  always  liable  to  be  precipitated  along  with 
the  persulphide  in  its  preparation,  and  also  to  the  instability  of  the 
compound,  its  exact  composition  has  been  the  subject  of  some 
doubt.  By  subjecting  the  crude  oil  as  first  precipitated  to  careful 
fractional  distillation  tinder  greatly  reduced  pressure,  Bloch  and 
Hohn  (Ber.,  1908)  have  not  only  succeeded  in  obtaining  the  disul- 
phide  in  a  sufficiently  pure  state  to  establish  its  formula,  but  have 
also  isolated  a  second  sulphide  having  the  composition  H2S3. 
This  compound,  hydrogen  trisulphide,  is  somewhat  denser,  specific 
gravity  1.496,  and  less  volatile  than  the  disulphide,  but  otherwise 
closely  resembles  it  in  properties. 


COMPOUNDS  OF  SULPHUR  WITH  CHLORINE. 

Two  of  these  compounds  exist  at  ordinary  temperatures,  while  a 
third  is  only  known  at  temperatures  below  -  22°. 

1.  Disulphur  dichloride  or  sulphothionyl  chloride     S2C12. 

2.  Sulphur  dichloride      ......     SC12. 

3.  Sulphur  tetrachloride         .....     SC14.  * 

Disulphur  Dichloride,  S2C12.—  This  substance  is  obtained  by 
passing  dry  chlorine  over  the  surface  of  heated  sulphur,  contained 
in  a  retort  ;  the  compound,  which  distils  away  as  fast  as  it  is 
formed,  condenses  in  the  receiver  as  a  yellow  liquid  — 

S2  +  C12  =  S2C12. 

Properties.  —  The  redistilled  liquid  is  an  amber-coloured  fuming 
substance  with  a  disagreeable  penetrating  odour,  the  vapour  of 
which  irritates  the  eyes.  Its  specific  gravity  is  1.709,  and  it  boils 
at  138.1°.  In  contact  with  water  it  gradually  decomposes  into 
hydrochloric  acid  and  sulphur  dioxide,  with  the  precipitation  of 
sulphur.  The  action  takes  place  in  two  stages,  thiosulphuric  acid 
being  formed  as  an  intermediate  product,  thus  — 


(a)  2 

03)      H2S203=H,S03-fS. 


414  Inorganic  Chemistry 

Disulphur  dichloride  dissolves  sulphur  with  great  readiness,  and 
the  solution  so  obtained  is  largely  employed  in  the  process  of 
vulcanising  indiarubber. 

This  compound  is  the  most  stable  of  the  three  chlorides  of 
sulphur.  From  the  fact  that  it  contains  chlorine  and  sulphur  in 
the  proportion  of  one  atom  of  each  element,  it  is  sometimes  called 
sulphur  monochloride  ;  but  as  its  vapour-density  (67.5)  shows  that 
it  contains  two  atoms  of  each  element  in  the  molecule,  the  use  of 
the  word  monochloride  is  calculated  to  mislead.  The  name  sul- 
phothionyl  chloride  indicates  its  analogy  to  thionyl  chloride,  SOC12> 
from  which  it  may  be  regarded  as  being  derived,  by  the  replace- 
ment of  the  oxygen  by  an  atom  of  sulphur,  O  =  S<^  ~,  ;  S  =  S<^  ~/ 

Sulphur  Dichloride,  SC12. — This  compound  is  obtained  by  passing  a  stream 
of  dry  chlorine  into  disulphur  dichloride  at  a  temperature  not  above  o°.  When 
the  maximum  amount  of  chlorine  is  absorbed,  the  liquid  assumes  a  dark 
reddish-brown  colour.  Excess  of  chlorine  is  removed  by  passing  a  stream  of 
carbon  dioxide  through  the  liquid. 

Sulphur  dichloride  rapidly  dissociates  with  rise  of  temperature  into  free 
chlorine  and  disulphur  dichloride;  at  +20°  this  decomposition  amounts  to 
6.5  per  cent.,  at  50°,  24.59  Per  cent.,  and  at  100°,  80.85  Per  cent-  On  boiling 
the  compound,  therefore,  chlorine  is  evolved,  and  the  disulphur  dichloride 
remains  behind. 

In  contact  with  water  it  is  decomposed  in  the  same  manner  as  the  more 
stable  compound. 

Sulphur  Tetrachloride,  SC14.— This  compound  only  exists  at  temperatures 
below  —  22°,  and  is  produced  by  saturating  sulphur  dichloride  with  chlorine  at 
that  temperature.  It  dissociates  very  rapidly  as  the  temperature  i/ses ;  thus, 
at  7°  above  the  temperature  at  which  it  is  formed,  viz.,  at  —15°,  this  decom- 
position amounts  to  58.05  per  cent.  At  -2°,  88.07  Per  cent-  °f  tne  compound 
dissociates,  while  at  +6.2°  the  percentage  rises  to  97.57. 

The  compound  is  decomposed  by  water  with  violence  into  sulphur  dioxide 
and  hydrochloric  acid — 

SC14+2H2O=SO2+4HC1. 

Compounds  of  Sulphur  with  Bromine  and  Iodine  have  been  obtained, 
corresponding  to  S2C12.  S2Br2  as  a  red-coloured  liquid,  boiling  with  partial 
decomposition  at  200°  ;  and  S2I2  as  a  dark-grey  crystalline  solid,  which  melts 
at  a  temperature  about  60°. 

OXIDES  AND   OXYACIDS   OF   SULPHUR. 
Four  oxides  of  sulphur  are  known,  namely — 

(i.)  Sulphur  sesquioxide  (hyposulphurous  anhydride)  S2O3. 

(2.)  Sulphur  dioxide  (sulphurous  anhydride)       .         .  SO2. 

(3.)  Sulphur  trioxide  (sulphuric  anhydride)         .         .  SO3. 

(4.)  Persulphuric  anhydride S2O7. 


Sulphur  Dioxide  415 

These  four  oxides  give  rise  respectively  to  the  acids,  hypo- 
sulphurous,  sulphurous,  sulphuric,  and  persulphuric,  besides 
which  several  other  sulphur  acids  are  known — 

HO'SO 

Hyposulphurous  acid      .         .     H2S2O4  HO-SO 

HO 

Sulphurous  acid       .         .         «     H2SO3  /SO 

HO 

HOv 

Sulphurrc  acid         .         .         .     H2SO4  /SO2 

HO' 

HO. 
Permonosulphuric  acid  .         .     H9SO5  /SO2 

HO-CK 
/OH     HO. 
Perdisulphuric  acid         .         .     H2S2O8    SO2\  /SO2 

S~\  T  T  T  T  /~\ 

Pyrosulphuric     acid     (Nord-    )  j_j  g  Q      SO— O ^SO2 

hausen  sulphuric  acid)      .  (      227 

HO. 
Thiosulphuric  acid  *        .         .     H2S2O3  >SO2 

HSX 

Besides  these  acids,  there  is  a  series  known  under  the  general 
name  of  the  polythionic  acids.  They  may  be  regarded  as  being 
derived  from  dithionic  acid,  which  is  the  first  of  the  series,  by  the 
absorption  into  the  molecule  of  various'  quantities  of  sulphur. 
Four  of  these  acids  are  believed  to  exist,  viz.  : — 

Dithionic  acid  (sometimes  called?  T_T  c  o  HO'SO2  ) 

hyposulphuric  acid)  f  n2^6  HO'SO2  J 

Trithionic  acid    ....     H2S3O6  O'SQ2  (  ^' 

2  j 

Tetrathionic  acid         .         .         .     H2S4O6  [O-SO    v  ^2> 

Pentathionic  acid        .         .         .     H2S5O6  Ho'sQ2  C  S3' 

2  * 

SULPHUR   DIOXIDE. 

Formula,  SO2.     Molecular  weight =64. 07.'    Density=32.o35. 

Occurrence. — This  compound  is  met  with  in  the  gaseous 
emanations  from  volcanoes,  and  in  solution  in  certain  volcanic 
springs.  It  is  also  present  in  the  air  of  towns,  being  derived 
from  the  combustion  of  the  sulphur  compounds  present  in  coal. 

*  This  acid  is  sometimes  incorrectly  called  hyposulphurous  acid,  its  sodium 
salt  being  known  as  sodium  hyposulphite:  the  so-called  "hypo"  of  the  photo- 
graphers. 


41  6  Inorganic  Chemistry 

Modes  Of  Formation.—  (  i.)  Sulphur  dioxide  is  formed  when 
sulphur  burns  in  air  or  oxygen  — 

S  +  O2  =  SO2. 

At  the  same  time  small  quantities  of  sulphur  trioxide  are  formed, 
which  render  the  gas  obtained  by  this  combustion  more  or  less 
foggy. 

(2.)  Sulphur  dioxide  may  also  be  obtained  by  heating  sulphur 
with  a  metallic  peroxide,  such  as  manganese  dioxide,  thus  — 

S2  +  MnO2  =  SO2+  MnS. 

(3.)  It  is  obtained  when  such  metallic  sulphides  as  copper  pyrites 
or  iron  pyrites  are  roasted  in  a  current  of  air,  the  metal  being  con- 
verted into  oxide,  thus  — 


(4.)  The  most  convenient  laboratory  process  for  preparing  sul- 
phur dioxide  consists  in  heating  sulphuric  acid  with  copper,  the 
final  products  of  the  reaction  being  copper  sulphate,  water,  and 
sulphur  dioxide  *  — 


The  metals  mercury  or  silver  may  be  substituted  for  copper,  but 
in  practice  the  latter  metal  is  usually  employed. 

(5.)  Sulphur  dioxide  is  also  formed  when  sulphuric  acid  is  heated 
with  sulphur,  the  oxidation  of  the  sulphur  and  the  reduction  of  the 
sulphuric  acid  going  on  simultaneously  — 


(6.)  The  reduction  of  sulphuric  acid  may  be  brought  about  by 
means  of  carbon  ;  thus,  if  sulphuric  acid  be  heated  with  carbon,  the 
latter  is  oxidised  to  carbon  dioxide,  and  the  acid  is  reduced  to 
sulphur  dioxide  — 


This  method  is  employed  on  a  large  scale  for  the  preparation  of 
alkaline  sulphides.  The  carbon  dioxide  which  accompanies  the 
sulphur  dioxide,  not  being  soluble  to  any  extent  in  water  containing 
sulphurous  acid,  is  not  in  any  way  detrimental. 

*  Secondary  reactions  go  on  simultaneously,  resulting  in  the  formation  of 
black  cuprous  sulphide  — 

5Cu  +  4H2S04  =  3CuSO4  +  4 


Sulphur  Dioxide  417 

(7.)  Sulphur  dioxide  is  formed  by  the  decomposition  of  a  sulphite 
by  dilute  sulphuric  acid,  thus — 

Na2SO3+  H2SO4=  Na2SO4  +  H2O  +  SO2. 

Properties. — Sulphur  dioxide  is  a  colourless  gas,  having  the 
well-known  suffocating  smell  usually  associated  with  burning 
sulphur.  The  gas  will  not  burn  in  the  air,  nor  will  it  support  the 
combustion  of  ordinary  combustibles  :  a  taper  introduced  into  the 
gas  is  instantly  extinguished.  Sulphur  dioxide  is  more  than  twice 
as  heavy  as  air,  its  specific  gravity  being  2.211  (air=i).  On  this 
account  it  is  readily  collected  by  displacement ;  it  cannot  be 
collected  over  water  on  account  of  its  solubility  in  that  liquid,  but 
may  be  collected  over  mercury.  The  solubility  of  sulphur  dioxide 
in  water  at  various  temperatures  is  seen  by  the  following  figures — 

i  vol.  of  water  at  o°  dissolves  79.789  vols.  SO2. 

»  »         20°         »         39-374 

„  „         40°         „         18.766        „ 

The  solution  is  strongly  acid,  and  is  regarded  as  sulphurous 
acid,  the  gas  having  entered  into  chemical  union  with  the  water — 

SO2+H2O  =  H2SO3. 

On  cooling  a  saturated  solution  of  sulphur  dioxide  to  o°,  a  solid 
crystalline  hydrate  is  deposited,  having  the  composition  H2SO3, 
8H2O.* 

When  the  solution  is  boiled  the  whole  of  the  sulphur  dioxide  is 
expelled. 

Sulphur  dioxide  is  an  easily  liquefied  gas.  At  o°  a  pressure  of 
1.53  atmospheres  is  sufficient  to  condense  it,  while  at  ordinary 
pressures  it  may  be  liquefied  by  a  cold  of  — 10°.  Its  critical 
temperature  is  155.4°. 

To  obtain  liquid  sulphur  dioxide,  the  gas,  as  evolved  from  the 
action  of  sulphuric  acid  upon  copper,  is  dried  by  being  passed 
through  a  bottle  containing  sulphuric  acid,  and  is  then  passed 
through  a  gas-condensing  tube  (Fig.  113)  immersed  in  a  freezing- 
mixture.  The  gas  at  once  condenses  in  the  bulb  of  the  apparatus 
as  a  colourless,  transparent,  mobile  liquid,  which  boils  at  —8°. 
When  the  liquid  is  cooled  to  —  76°  it  solidifies  to  a  transparent,  ice- 
like  mass. 

*  Several  hydrates  of  sulphurous  acid  have  been  obtained,  H2SO3,6H2O; 
H2SO310H2O;  HaSO3,14Hj,'J. 

2  D 


Inorganic  Chemistry 

Liquid  sulphur  dioxide  is  largely  employed  as  a  refrigerating 
agent,  low  temperatures  being  obtained  by  its  rapid  evaporation 
under  reduced  pressure.  The  liquid  dissolves  phosphorus,  iodine, 
sulphur,  and  many  resins.  When  thrown  upon  water  a  portion  of 
the  liquid  dissolves,  and  owing  to  the  reduction  of  temperature 
caused  by  the  rapid  evaporation  of  the  remainder,  a  quantity  of 
the  water  is  frozen.  The  ice  so  produced  contains  a  large  pro- 
portion of  the  solid  hydrate,  H2SO3,8H2O. 

Although  sulphur  dioxide  is  incapable  of  supporting  the  com- 


FIG.  113. 

bustion  of  ordinary  combustibles,  many  metals  will  take  fire  and 
burn  when  heated  in  the  gas.  Thus,  when  finely  divided  iron  is 
heated  in  a  stream  of  sulphur  dioxide  it  burns,  forming  sulphide 
and  oxide  of  the  metal. 

It  also  unites  with  many  metallic  peroxides,  and  often  with  so 
much  energy  as  to  give  rise  to  light  and  heat.  Thus,  when  passed 
over  peroxide  of  lead,  the  mass  glows  spontaneously  in  the  gas, 
and  lead  sulphate  is  produced — 

+  S02=PbS04. 


Sulphur  Dioxide  419 

Or  if  sodium  peroxide  is  dusted  into  a  cylinder  of  the  gas,  the 
peroxide  burns  with  a  brilliant  light,  yielding  sodium  sulphate  — 
Na2O2+SO2  =  Na2SO4. 

Sulphur  dioxide  is  decomposed  by  the  influence  of  strong  light. 
If  a  concentrated  beam  of  electric  light  be  passed  through  a  vessel 
filled  with  gaseous  sulphur  dioxide,  the  gas  at  first  will  appear 
perfectly  transparent  and  clear  ;  but  in  the  course  of  a  few  minutes 
the  track  of  the  beam  will  become  more  and  more  visible  as  it 
traverses  the  gas,  owing  to  the  formation  of  thin  clouds  of  sulphur 
trioxide  and  sulphur,  until  the  atmosphere  of  the  vessel  appears 
to  be  filled  with  fog  (Fig.  114)  — 


After  the  lapse  of  a  short  time,  if  the  vessel  be  removed  from 


FIG.  114. 

the  strong  light,  the  atmosphere  will  once  more  become  clear, 
owing  to  the  reformation  of  sulphur  dioxide. 

Sulphur  dioxide  possesses  powerful  bleaching  properties  when 
in  the  presence  of  water.  Its  bleaching  action  is  due  to  its 
absorption  of  oxygen  from  water,  and  consequent  liberation  of 
hydrogen,  thus  — 


The  hydrogen  so  set  free  reduces  the  colouring-matter,  with  the 
formation  of  colourless  compounds  :  the  action  in  this  case  being 
the  reverse  to  that  which  takes  place  with  chlorine.  In  some 
instances,  the  bleaching  is  due  to  the  formation  of  a  colourless 
compound,  by  the  direct  combination  of  sulphur  dioxide  with  the 
colouring-matter,  as  the  original  colour  may  often  be  restored  by 
treatment  with  dilute  sulphuric  acid,  or  by  weak  alkaline  solutions. 


420  Inorganic  Chemistry 

Thus,  by  passing  sulphur  dioxide  into  an  infusion  of  rose  leaves, 
the  red  colour  of  the  liquid  is  quickly  discharged,  but  on  the  addi- 
tion of  a  small  quantity  of  sulphuric  acid  the  colour  is  restored. 

Sulphur  dioxide  is  employed  in  bleaching  materials  that  would 
be  injured  by  exposure  to  chlorine,  such  as  straw,  silk,  wool, 
sponge,  &c.,  and  the  familiar  yellow  colour  which  gradually  comes 
over  a  sponge  or  a  piece  of  bleached  flannel  when  it  is  washed  with 
soap  is  an  illustration  of  the  power  of  alkalies  to  restore  the  original 
colour  to  materials  that  have  been  bleached  by  this  substance. 

In  the  presence  of  water  sulphur  dioxide  converts  chlorine  into 
hydrochloric  acid,  and  on  this  account  is  employed  as  an  "  anti- 
chlor"  — 


In  the  same  way  it  acts  upon  iodine,  with  the  formation  of 
hydriodic  acid  — 

SO2  +  2H2O  +  I2  =  2HI  +  H2SO4. 

In  the  case  of  iodine_,  however,  this  reaction  only  takes  place 
when  a  certain  degree  of  dilution  is  maintained,  for  in  a  more 
concentrated  condition  sulphuric  acid  is  reduced  by  hydriodic  acid 
into  sulphur  dioxide,  according  to  the  reverse  equation,  thus  — 
H2SO4  +  2HI  =  I2  +  2H,O  +  SO2. 

It  has  been  shown*  that  aqueous  sulphurous  acid  can  only  be 
completely  oxidised  by  iodine,  as  indicated  in  the  former  equation, 
when  the  proportion  of  sulphur  dioxide  does  not  exceed  0.05  per 
cent.  ;  when  the  amount  exceeds  this  proportion  the  second  reaction 
comes  into  operation. 

Sulphur  dioxide  brought  into  contact  with  iodic  acid,  or  an 
iodate,  is  oxidised  into  sulphuric  acid  and  liberates  iodine,  thus  — 


This  reaction  is  made  use  of  as  a  method  for  the  detection  of 
the  presence  of  sulphur  dioxide.  Paper  which  has  been  moistened 
with  a  solution  of  potassium  iodate  and  starch,  on  exposure  to 
sulphur  dioxide  is  at  once  turned  blue,  owing  to  the  liberated 
iodine  combining  with  the  starch. 

The  composition  of  sulphur  dioxide  may  be  determined  by  the 
combustion  of  sulphur  in  a  measured  volume  of  oxygen,  in  the 
apparatus  employed  for  showing  the  volume  composition  of  carbon 
dioxide  (Fig.  67).  After  the  fragment  of  sulphur  has  burnt,  and 
the  apparatus  has  been  allowed  to  cool,  it  will  be  seen  that  there 
is  no  alteration  in  the  volume  of  the  contained  gas  —  the  sulphur 

*  Bunsen. 


Sulphur  Trioxide  421 

dioxide  produced  occupying  the  same  volume  as  the  oxygen  used 
in  its  formation.  Sulphur  dioxide,  in  other  words,  contains  its 
own  volume  of  oxygen.  One  molecule,  therefore,  of  sulphur 
dioxide  contains  one  molecule  of  oxygen,  weighing  32.  But  the 
molecular  weight  of  sulphur  dioxide  is  64.06  ;  therefore  64.06  -  32  = 
32.06  =  the  weight  of  sulphur  contained  in  the  molecule  of  the  gas. 
Sulphur  dioxide,  therefore,  contains  in  the  molecule  one  atom  of 
sulphur  combined  with  two  atoms  of  oxygen,  hence  its  composition 
is  expressed  by  the  formula  SO2. 

Sulphurous  Acid  and  Sulphites.—  Sulphurous  acid  is  only 
known  in  solution  and  in  its  hydrates.  The  solution  smells 
of  sulphur  dioxide,  and  gradually  undergoes  decomposition  by 
absorption  of  oxygen.  The  acid  is  dibasic,  having  two  atoms  of 
hydrogen  replaceable  by  metals  ;  it  is  therefore  capable  of  form- 
ing two  series  of  salts,  according  to  whether  one  or  both  of  the 
hydrogen  atoms  are  replaced.  Thus,  by  its  action  upon  potassium 
hydroxide,  when  the  acid  is  in  excess,  the  so-called  acid  potassium 
sulphite,  or  hydrogen  potassium  sulphite,  is  obtained  — 
KHO  +  H2SO3=  H,0  +  HKSO3. 

Whereas,  if  the  metallic  hydroxide  be  in   excess,  the  normal 
potassium  sulphite  is  formed  — 


The  alkaline  sulphites  are  readily  soluble  in  water,   all   other 
normal  sulphites  being  either  difficult  of  solution  or  insoluble. 

SULPHUR  TRIOXIDE. 

Formula,  SO3.     Molecular  weight  =  80.  07.     Vapour  density  =40.  03. 
Modes  of  Formation.—  (  i.)  This  compound  is  produced  when 
a  mixture  of  sulphur  dioxide  and  oxygen  is  passed  over  heated 
spongy  platinum  or  platinised  asbestos  — 
2SO2+Oa  =  2SO3. 

On  leading  the  product  through  a  well-cooled  receiver,  the  sulphur 
trioxide  condenses  in  white  silky  needles.  This  method  has  been 
successfully  employed  on  a  commercial  scale.  The  mixture  of 
sulphur  dioxide  and  oxygen  is  obtained  by  allowing  ordinary  strong 
sulphuric  acid  to  drop  into  earthenware  retorts  heated  to  bright 
redness,  whereby  it  is  almost  entirely  broken  up  into  these  two 
gases  and  water,  thus  — 

H2SO4=SO2  +  O  +  H2O. 

The  gases  are  then  deprived  of  the  water,  by  passage  first  through 
a  condenser,  and  then  through  a  leaden  tower  containing  coke 


422  Inorganic  Chemistry 

moistened  with  sulphuric  acid,  and  are  finally  passed  over  heated 
platinised  asbestos  contained  in  glazed  earthenware  pipes. 

(2.)  Sulphur   trioxide   is  most  conveniently  obtained  by  gently 
heating  pyrosulphuric  acid  in  a  glass  retort.     The  trioxide  distils 
over  and  may  be  collected  in  a  well-cooled  receiver  — 
H2S,Or-H,S04  +  SO3. 

(3.)  It  may  also  be  obtained  by  heating  sodium  pyrosulphate  to 
bright  redness  — 


The  sodium  pyrosulphate  is  produced  when  hydrogen  sodium  sul- 
phate (so-called  bisulphatc  of  soda)  is  heated  to  about  300°,  thus  — 

2HNaSO4  =  H,0  +  Na,S,Or. 

And  on  account  of  this  origin  it  is  sometimes  termed  anhydrous 
sodium  bisulphate. 

(4.)  Sulphur  trioxide  can  also  be  produced  by  the  action  oi 
phosphorus  pentoxide  upon  sulphuric  acid.  This  most  powerful 
dehydrating  substance  withdraws  from  the  sulphuric  acid  the  ele- 
ments of  water  when  gently  heated,  thus  — 


The  trioxide  is  distilled  from  the  mixture,  and  the  metaphosphonc 
acid  remains  in  the  retort. 

Properties.—  Sulphur  trioxide  is  a  white,  silky-looking,  crystal- 
line substance,  which  melts  at  14.8°  and  boils  at  46°.  It  is  very 
volatile,  and  gives  off  dense  white  fumes  in  contact,  with  air,  owing 
to  the  combination  of  its  vapour  with  atmospheric;  moisture  to  form 
sulphuric  acid.  It  combines  with  water  with  great  energy  to  form 
sulphuric  acid  ;  a  fragment  of  the  compound  dropped  into  water 
dissolves  with  a  hissing  sound  resembling  the  quenching  of  red-hot 
iron  — 

SO3  +  H,,O  =  H2SO4. 

When  brought  in  contact  with  the  skin,  or  other  organic  matter 
containing  hydrogen  and  oxygen,  it  abstracts  these  elements  and 
produces  a  burnt  or  charred  effect  upon  the  substance.  Sulphur 
trioxide  unites  directly  with  barium  oxide,  BaO,  and  if  the  baryta 
be  dry  the  mass  becomes  incandescent,  owing  to  the  heat  of  the 
union,  and  barium  sulphate  is  formed  — 

BaO  +  SO3=BaSO4. 

When  the  vapour  of  sulphur  trio^ide  is  passed  through  a  red-hot 
tube,  it  is  broken  down  into  sulphur  dioxide  and  oxygen. 

When  the  trioxide  is  heated,  it  melts  to  a  colourless  liquid,  which 


Sulphur  Trioxide  423 

exhibits  a  remarkably  high  rate  of  expansion  by  heat ;  between 
25°  and  45°  its  mean  coefficient  of  expansion  is  0.0027,  nearly  three- 
fourths  of  the  expansion  coefficient  of  a  gas. 

Sulphur  Sesquioxide,  S2O^. — A  solution  of  this  compound  in  fuming 
sulphuric  acid  was  obtained  early  in  the  century  by  heating  flowers  of  sulphur 
with  Nordhausen  sulphuric  acid,  whereby  a  blue  .solution  was  obtained.  The 
substance  may  be  prepared  by  the  gradual  addition  of  dry  flowers  of  sulphur 
to  melted  sulphur  trioxido,  at  a  temperature  just  above  its  melting-point,  when 
a  malachite-green  crystalline  solid  separates  out. 

The  compound  is'  unstable  at  ordinary  temperatures,  being  resolved  into 
sulphur  dioxide  and  sulphur,  the  decomposition  taking  place  rapidly  upon 
gently  warming — 

2S2O3=S  +  3S02. 

If  the  sesquioxide  be  sealed  up  in  a  bent  glass  tube  and  gently  warmed,  the 
sulphur  dioxide  may  be  obtained  liquid  in  one  limb  of  the  tube. 

Hyposulphurous  Acid,*  H-jboOi. — This  compound  was  discovered  by 
Schutzenberger,  who  gave  to  it  The  formula  H2SO2.  The  later  investigations 
of  Bernsthen  prove  that  its  composition  is  expressed  by  the  formula  H2S2O4. 
It  is  obtained  by  the  reduction  of  sulphurous  acid  by  means  of  nascent 
hydrogen.  Thus,  when  zinc  is  acted  upon  by  an  aqueous  solution  of  sul- 
phurous acid,  no  hydrogen  is  evolved,  as  the  nascent  gas  combines  with 
oxygen  of  the  acid  to  form  water — • 

2H2S03  +  2H  =  2  H20  +  H2S,O4. 

The  solution  so  obtained  has  a  yellowish  colour,  and  possesses  powerful 
reducing  and  bleaching  properties. 

Sodium  hyposulphite  (Na2S.2O4)  may  be  obtained  by  the  action  of  zinc  upon 
a  cooled  concentrated  solution  of  hydrogen  sodium  sulphite  (HNaSO3),  air 
being  carefully  excluded,  a  double  sulphite  of  sodium  and  zinc  being  at  the 
same  time  produced,  thus — 

4H  XaSO3  +  Zn  =  Na.2S2O4  +  NagSOo,  ZnSO3  +  2H2O. 

The  greater  part  of  the  double  sodium-zinc  sulphite  is  deposited  as  crystals, 
the  rest  is  removed  by  adding  to  the  mother-liquor  about  four  times  its  volume 
of  alcohol,  in  a  closed  flask.  The  double  salt,  being  less  soluble  in  alcohol 
than  the  hyposulphite,  is  first  deposited,  and  the  clear  liquid  after  being  poured 
off  and  corked  up,  is  cooled,  when  it  solidifies  to  a  mass  of  crystals  of  nearly 
pure  sodium  hyposulphite.  The  crystals  in  the  wet  condition  are  rapidly 
oxidised  on  exposure  to  air  ;  but  if  quickly  pressed  between  blotting-paper  and 
dried  in  vacuo,  the  dry  salt  is  not  acted  upon  by  atmospheric  oxygen. 

The  acid  is  obtained  from  the  sodium  salt  by  the  action  upon  it  of  oxalic  acid. 

Sodium  hyposulphite  is  also  formed  when  a  solution  of  hydrogen  sodium 
sulphite  is  subjected  to  electrolysis,  the  nascent  hydrogen  developed  at  the 
negative  electrode  reducing  the  sulphite  to  hyposulphite  by  the  abstraction  of 
one  atom  of  oxygen. 

Properties. — Hyposulphurous  acid  is  an  extremely  unstable,  yellow-coloured 
liquid  which  rapidly  decomposes  into  sulphur  dioxide,  water,  and  sulphur — 

2H2S.2O4=3SO2+2H2O  +  S, 

a  certain  quantity  of  thiosulphuric  acid  being  formed  as  an  intermediate 
product  at  the  same  time— 

2H2S2O4 = H3S203  -f  2SO2  +  H2O. 


*  This  compound  (sometimes  called  kydnsvlphvrous  acid)  must  not  be  con- 
founded with  thiosulphuric  acid,  which  is  often  incorrectly  called  hypo- 
sulphurous  acid  (page  435). 


424  Inorganic  Chemistry 

The  acid  reduces  salts  of  silver  or  mercury,  with  precipitation  of  the  metal, 
thus  — 


The  sodium  salt  possesses  the  same  bleaching  and  reducing  powers  as  the 
acid,  and  when  wet  or  in  solution  it  rapidly  absorbs  oxygen  from  the  air  and 
is  converted  into  a  compound  known  as  sodium  metabisulphite  — 


The  relation  in  which  these  two  compounds  stand  to  each  other  will  perhaps 
be  more  evident  if  their  formulas  are  written  thus  — 
NaO-SO)    ,  n_NaO-SO)n 
NaO-SO  j  +U~  NaO-SO  /Ul 

Persulphuric  Anhydride,  S2O7,  is  formed  when  a  mixture  of  sulphur 
trioxide  and  oxygen  is  subjected  to  the  prolonged  action  of  the  silent  electric 
discharge.  The  compound  is  very  unstable,  breaking  up  even  at  compara- 
tively low  temperatures  into  its  generators. 

Perdisulphuric    Acid   (Persulphuric    Acid],    H2S0O8.  —  When    sulphuric 

+ 
acid  is   diluted  with  water  it  not  only  dissociates  into  hydrion  H  and  sul- 

phanion  SO4,  but  also  into  hydrion  and  hydrosulphanion  HSO4,  the  extent  to 
which  dissociation  in  these  two  directions  takes  place  depending  upon  the 
degree  of  dilution.  Within  certain  limits  the  less  the  acid  is  diluted  the 

larger  is  the  proportion  of  HSO4  ions  present 

When,  therefore,  moderately  strong  sulphuric  acid  is  electrolysed,  the  HSO4 
ions  on  discharging  at  the  anode  unite  to  form  molecules  of  perdisulphuric 
acid  (HSO4)2  or  H.,S2O8.  The  pure  acid  has  not  been  isolated,  its  aqueous 
solutions  rapidly  decomposing  into  sulphuric  acid  and  oxygen.  The  acid  is 
dibasic,  and  its  salts  (usually  called  persulphates,  but  more  correctly  per- 
disulphates)  have  the  general  formula  M'2S.2O8. 

The  potassium  salt  is  obtained  by  the  electrolysis  of  a  strong  solution  of 

potassium  hydrogen  sulphate,  which  will  contain  the  ions  K,  HSO4.  The 
acid  which  is  first  formed  by  the  union  of  the  discharged  negative  ions  at  the 
anode,  there  interacts  with  the  potassium  hydrogen  sulphate  forming  the 
sparingly  soluble  persulphate  which  separates  out.  The  salt  may  be  freed 
from  the  acid  sulphate  by  recrystallisation. 

Barium  perdisulphate,  BaS2Og,  4H2O,  is  soluble  in  water  ;  100  parts  of 
water  at  o°  dissolving  52-2  parts  of  the  salt.  On  this  account  barium  chloride 
gives  no  precipitate  with  a  solution  of  a  persulphate,  which  distinguishes  these 
salts  from  the  sulphates.  On  warming,  however,  the  persulphate  is  decom- 
posed into  sulphate  and  oxygen  — 

BaS208  +  H,0  =  BaS04  +  H2SO4  +  O  . 

In  the  solid  state  the  persulphates  are  stable  salts,  but  their  aqueous  solu- 
tions gradually  evolve  oxygen  and  pass  into  sulphates.  The  reactions  of  both 
the  acid  and  its  salts  are  therefore  those  of  strong  oxidising  agents. 

Permonosulplmric  Acid  (Card's  Acid],  H2SO5,  is  obtained  when  potassium 


Sulphuric  Acid  425 

perdisulphate  is  treated  with  strong  sulphuric  acid  and  the  mixture  poured 
upon  crushed  ice,  both  operations  being  conducted  in  a  freezing  mixture — 
K2S2O8-f  H0SO4=  K2SO4+  H2S2O8. 
H2S2O8+  H2O=  H2SO4  +  H2SO5. 

The  relation  in  which  these  two  persulphuric  acids  stand  to  each  other  will 
be  seen  by  a  consideration  of  their  constitutional  formulas — 

/OH  /OH     H0\ 

SO.<  ;       SOa<T  )SOo. 

-\O-OH  >O O/ 

Permonosulphuric  acid  is  regarded  as  a  monobasic  Stid,  the  "acidic"  OH 
group  being  the  one  directly  associated  with  the  sulphur.  No  pure  salts  have 
yet  been  obtained. 

SULPHURIC  ACID. 

Formula,  H2SO4. 

Modes  of  Formation. — (i.)  This  acid  is  formed  when  sulphur 
trioxide  is  dissolved  in  water — 

SO3+H,O  =  H2SO4. 

(2.)  It  is  also  formed  by  the  direct  union  of  sulphur  dioxide  with 
hydrogen  peroxide — • 

SO.2-fH2O2=H,SO4. 

(3.)  An  aqueous  solution  of  sulphur  dioxide  gradually  absorbs 
oxygen,  and  is  converted  into  sulphuric  acid — 
H2SO3+O  =  H2SO4. 

(4.)  Manufacture  Of  Sulphuric  Acid.— Sulphur  dioxide  is  un- 
able to  absorb  an  additional  atom  of  oxygen,  and  so  pass  into 
sulphur  trioxide,  without  the  aid  of  some  third  substance  which 
can  act  as  a  catalytic  agent  or  a  carrier  of  oxygen.  The  material 
which  is  employed  for  this  purpose  in  the  process  by  which  sul- 
phuric acid  is  manufactured  is  one  of  the  oxides  of  nitrogen,  which 
is  capable  of  giving  up  oxygen  to  the  sulphur  dioxide,  and  of  again 
taking  up  oxygen  from  the  air.  Thus,  nitrogen  peroxide  (NO.,), 
by  the  loss  of  one  atom  of  oxygen,  is  reduced  to  nitric  oxide,  NO  ; 
which  in  its  turn  combines  with  atmospheric  oxygen  and  is  re- 
converted into  nitrogen  peroxide.  Therefore,  when  sulphur  dioxide 
and  oxygen  are  mixed  with  nitrogen  peroxide  in  the  presence  of 
steam,  a  series  of  reactions  takes  place,  the  final  result  of  which 
is  that  the  oxygen  is  caused  to  combine  with  the  sulphur  dioxide 
and  water,  with  the  formation  of  sulphuric  acid — 
S02  +  0  +  H20  =  H2S04. 

The  nitrogen  peroxide  at  the  end  of  the  reaction  is  unchanged, 
and  is  able  to  react  in  the  same  series  of  changes  over  and  over 
again,  thus  transforming,  theoretically,  an  unlimited,  and,  in 


426  Inorganic  Chemistry 

practice,  a  relatively  large  quantity  of  sulphur  dioxide  into  sul- 
phuric acid. 

The  series  of  changes  that  gives  rise  to  the  ultimate  product  is 
the  following  :  —  The  sulphur  dioxide,  nitrogen  peroxide,  and  water 
give  rise,  in  the  first  place,  to  the  formation  of  nitro-sulphonic  acid 
and  a  molecule  of  nitric  oxide  — 


(i.)  2S02  +  3N02  +  H20  =  2H(NO)S04 

Nitro-sulphonic  acid  (sometimes  called  nitro-sulphuric  acid,  and 
nitrosyl  sulphate}  may  be  regarded  as  sulphuric  acid  in  which  one 
of  the  hydrogen  atoms  is  replaced  by  the  group  (NO),  thus, 

/OH 
SO2<f  p.     .,  ~    in  which  case  the  nitrogen  is  linked  to  the  sulphur 

by  the  intervention  of  oxygen  ;  or  it  may  be  considered  as  derived 
from  sulphuric  acid  by  the  replacement  of  one  of  the  groups  (HO) 

/OH 
by  the  group  NO2,  SO2^  N  ~      when    the    nitrogen    is    directly 

attached  to  the  sulphur.  The  substance  is  a  white  crystalline 
compound  which  in  the  presence  of  water  is  instantly  decomposed 
into  sulphuric  acid  and  a  mixture  of  nitric  oxide  and  nitrogen 
peroxide,  thus  — 

(2.)  2SO2(HO)(NO2)  +  H2O  =  2H2SO4  +  NO  4-  NO2. 

The  nitric  oxide  in  this  and  the  former  reaction,  on  coming  in 
contact  with  the  atmospheric  oxygen,  is  at  once  reconverted  into 
nitrogen  peroxide  — 

(3.)  NO  +  0  =  N02. 

In  the  process  of  the  manufacture  the  crystalline  compound 
SO2(HO)(NO2)  (known  as  chamber  crystals]  is  not  actually  isolated, 
unless  from  accidental  causes  the  supply  of  water  is  in  deficit,  the 
production  of  these  crystals  being  regarded  as  an  indication  that 
the  process  is  not  being  well  carried  out. 

The  formation  of  sulphuric  acid  by  these  reactions,  with  the 
intermediate  production  of  the  chamber  crystals,  may  be  carried 
out  on  a  small  scale  by  means  of  the  apparatus  shown  in  Fig.  115. 
A  large  flask,  F,  is  fitted  with  a  cork,  through  which  pass  five 
tubes  :  three  of  these  are  connected  to  separate  two-necked  bottles 
containing  sulphuric  acid,  through  which  can  be  delivered  respec- 
tively, nitric  oxide,  sulphur  dioxide,  and  oxygen.  The  fourth  tube 
is  attached  to  a  flask  in  which  water  may  be  boiled,  and  through 
which  oxygen  can  be  passed,  and  the  fifth  tube  (not  shown  in 


Sulphuric  Acid 


427 


the  figure)  serves  as  an  exit.  A  quantity  of  oxygen  is  first  passed 
into  the  large  flask  through  the  drying-bottle  D,  and  sufficient 
nitric  oxide  is  then  allowed  to  enter,  to  form  deep  red  vapours  ;  at 
the  same  time  sulphur  dioxide  is  passed  in  through  the  bottle  S. 
In  order  to  introduce  a  small  quantity  of  moisture,  oxygen  is 
allowed  to  enter  through  the  flask  of  boiling  water,  and  in  a  few 
moments  large  white  crystals  begin  to  form  all  over  the  interior 
of  the  flask,  and  rapidly  spread  until  the  whole  surface  is 
covered. 

In  order  to  show  the  second  reaction  in  the  cycle,  the  gaseous 


FIG.  115. 

contents  of  the  flask  may  be  swept  out  by  means  of  a  rapid  stream 
of  oxygen  passed  in  through  the  drying-bottle  D  ;  and  when  the 
atmosphere  within  the  apparatus  is  colourless,  a  quantity  of  steam 
is  driven  in  from  the  small  flask.  The  chamber  crystals  will 
be  seen  to  dissolve  with  effervescence,  and  the  flask  once  more 
becomes  filled  with  brown  fumes.  The  nitric  oxide  evolved  by 
the  decomposition  of  the  nitrosyl  sulphate,  coming  in  contact  with 
the  oxygen  within  the  flask,  at  once  regenerates  nitrogen  peroxide, 
in  accordance  with  equation  No.  3. 

The  solution  formed  in  the  flask  will  be  found  to  yield  a  pre- 
cipitate of  barium  sulphate,  on  the  addition  to  it  of  a  soluble 
barium  salt. 

On  a  manufacturing  scale,  the  combination  of  the  reacting  gases 
and  vapours  which  gives  rise  to  the  sulphuric  acid  takes  place  in 


428  Inorganic  Chemistry 

large  leaden  chambers,  usually  about  100  feet  long,  25  feet  wide, 
and  20  feet  high,  having  therefore  a  capacity  of  50,000  cubic  feet, 
several  of  such  chambers  being  placed  in  series.  Into  these 
chambers  there  is  delivered  sulphur  dioxide,  air,  oxides  of  nitrogen, 
and  steam. 

The  plant  employed  for  the  manufacture  of  sulphuric  acid  con- 
sists broadly  of  four  parts,  i.  Apparatus  for  generating  sulphur 
dioxide.  2.  Apparatus  for  producing  oxides  of  nitrogen.  3.  Appa- 
ratus for  absorbing  oxides  of  nitrogen  from  the  gases  leaving  the 
chambers.  4.  The  chambers  in  which  the  reactions  are  made. 

(i.)  Pyrites  Burners.  —  The  sulphur  dioxide  is  obtained  either  by 
burning  native  sulphur,  or  roasting  the  "  spent  oxide  "  of  the  gas 
works  (see  Sulphur),  or  by  roasting  pyrites,  the  latter  being  the 
most  general  method.  The  pyrites  burner  (Fig.  116,  B)  is  essen- 
tially a  small  furnace  or  kiln  in  which  the  ore  is  heated,  and  in 
which  the  admission  of  air  can  be  duly  regulated,  as  not  only  is  it 
necessary  to  admit  sufficient  air  to  completely  burn  the  whole  of 
the  sulphur,  and  so  prevent  any  volatilisation  of  it  in  an  unburnt 
condition,  but  also  to  supply  the  requisite  volume  of  oxygen  for 
the  requirements  of  the  reactions  which  are  to  go  on  within  the 
chamber.  Too  large  a  volume  of  air  must  be  avoided,  in  order  not 
to  unduly  dilute  the  chamber  gases. 

(2.)  If  no  loss  of  nitrogen  peroxide  took  place  during  the  cycle 
of  changes,  the  same  quantity  of  this  gas  would  convert  an  infinite 
amount  of  sulphur  dioxide  and  water  into  sulphuric  acid,  but  in 
practice,  owing  to  leakage,  defective  absorption,  and  the  reduction 
of  a  certain  percentage  of  this  compound  into  nitrous  oxide,  it  is 
necessary  to  constantly  replenish  the  supply.  This  is  usually  done 
by  generating  a  small  quantity  of  nitric  acid  (by  the  action  of 
sulphuric  acid  upon  nitre)  in  earthenware  pots,  which  are  usually 
placed  in  an  enlarged  part  of  the  flue  of  the  pyrites  burner,  known 
as  the  "nitre  oven,"  and  which  is  provided  with  a  door  for  the 
introduction  of  the  pots  (Fig.  116,  N).  The  heated  gases  playing 
upon  these  pots  promotes  the  evolution  of  the  nitric  acid,  which  in 
contact  with  sulphur  dioxide  is  at  once  decomposed  according  to 
the  equation  — 

2HNO3+SO2=H2SO4 


It  is  found  that  to  make  up  for  the  loss  of  nitrogen  peroxide, 
about  three  to  four  parts  of  nitre  are  required  for  every  100  parts 
of  sulphur  burnt  as  pyrites. 


Sulphuric  A  cid  429 

(3.)  The  apparatus  for  the  absorption  of  the  nitrogen  peroxide 
from  the  gases  that  are  drawn  from  the  chamber  at  the  end  of  the 
series  is  known  as  the  "  Gay-Lussac  Tower"  (Fig.  116,  T).  This 
consists  of  a  square  leaden  tower  filled  with  fragments  of  coke, 
and  down  which  there  is  caused  to  slowly  percolate  a  stream  of 
cold  strong  sulphuric  acid,  the  acid  being  evenly  spread  over  the 
mass  of  coke  by  a  special  distributing  arrangement.  The  nitrogen 
peroxide  is  absorbed  by  the  acid,  with  the  formation  of  nitro- 
sulphonic  acid,  SO2(HO)(NO.>).  In  order  to  make  use  of  the 
absorbed  nitroxygen  compound,  the  acid  which  flows  from  the 
Gay-Lussac  tower  is  pumped  to  the  top  of  another  very  similar 
tower,  situated  between  the  "  burners  "  and  the  first  of  the  cham- 
bers, and  known  as  the  "  Glover  Tower,"  G.  The  hot  gases  from 
the  burners,  consisting  of  sulphur  dioxide,  nitrogen,  and  oxygen, 
together  with  the  small  quantities  of  nitrogen  peroxide  from  the 
nitre  pots,  are  made  to  pass  up  this  tower  on  their  way  to  the  first 
chamber,  and  meeting  with  the  descending  stream  of  nitro-sul- 
phonic  acid  as  it  runs  over  the  bricks  or  flints  with  which  the  tower 
is  filled,  denitrification  of  the  latter  takes  place,  thus  — 


2SO.,(HO)(NOjj)  -1-  SO2  +  2H,O  =  2NO  +  3SO.,(HO)(HO), 
~or3H~.,SO4. 

The  nitric  oxide  thus  evolved,  in  presence  of  the  atmospheric 
oxygen,  is  converted  into  nitrogen  peroxide,  and  swept  along  with 
the  other  gases  into  the  chambers. 

In  practice  it  is  usual  to  deliver  down  the  Glover  tower,  besides 
the  nitro-sulphonic  acid,  a  quantity  of  "chamber  acid"  from  a 
separate  tank.  The  effect  of  the  heated  gases  upon  this  dilute 
acid  is  to  remove  a  portion  of  the  water  from  it,  thereby  effecting 
its  partial  concentration,  and  furnishing  the  water  demanded  by 
the  above  equation.  It  will  be  seen,  therefore,  that  there  is  a 
scrubber  tower  at  each  end  of  the  series  of  chambers,  the 
"Gay-Lussac"  at  the  exit,  where  nitrogen  peroxide  is  absorbed; 
and  the  "Glover"  at  the  commencement,  where  the  dissolved 
nitrogen  compound  is  again  liberated  and  returned  to  the 
chambers. 

(4.)  The  chambers  are  made  of  sheet  lead,  connected  togethei 
by  fusing  the  edges  by  means  of  an  oxyhydrogen  flame,  without 
the  intervention  of  solder,  as  the  presence  of  another  metal  gives 
rise  to  the  rapid  corrosion  of  the  lead  on  account  of  galvanic 


43° 


Inorganic  Chemistry 


action  being  set  up  ;  this  method  of  joining  the  lead  is  known  as 
autogenous  soldering.  The  enormous  leaden  chamber  is  supported 
in  a  framework  of  wood,  to  which  the  lead  is  secured  by  bands  of 


FIG.  116. 


B. —Double  row  of  pyrites  burners,  placed 
back  to  back :  one  being  shown 
open. 

N. — Hearth  where  the  nitre  pots  are 
placed  :  one  shown  as  open. 

G. — Glover  Tower,  with  two  tanks  at  top: 
one  for  the  nitro-sulphuric  acid  de- 
rived from  the  Gay-Lussac  tower, 
the  other  for  the  "chamber  acid." 


These  acids  are  forced  up  from  the 

leaden  vessels  E,  called  "  eggs." 
C. — Leaden  chamber,  of  which  there  are 

three  shown  in  the  figure. 
P. — Pipe  conveying   the   gases  from   the 

third   chamber   to  the    Gay-Lussac 

tower. 
T.—  Gay-Lussac  Tower.     The  tanks  at  tl  e 

top  of  this    and    the  Glover  tower 

are  enclosed  in  wooden  sheds. 


the  same  metal,  and  the  whole  is  usually  supported  on  iron  or 
brick  pillars. 

The  general  arrangement  of  a  modern  sulphuric  acid  works  is 


Sulphuric  Acid  431 

seen  in  Fig.  116.  The  gases  from  the  double  row  of  pyrites 
burners  B  are  led  through  the  Glover  tower  G,  where  they 
effect  the  denitrification  of  the  nitro-sulphonic  acid,  as  already 
explained.  From  this  tower  they  are  delivered  into  the  series  of 
chambers,  where  they  meet  with  the  necessary  supply  of  steam. 
The  acid  collects  upon  the  floor  of  the  chambers,  and  samples  are 
constantly  drawn  off  by  means  of  an  arrangement  known  as  a 
drip-pipe,  which,  acting  in  a  manner  similar  to  a  rain  gauge,  indi- 
cates the  progress  of  the  processes  going  on  within.  The  gases, 
after  being  drawn  through  the  entire  series  of  chambers  by  means 
of  the  draught  caused  by  the  tall  chimney,  are  finally  passed  up 
the  Gay-Lussac  tower  T,  where  all  the  nitrogen  peroxide  is  ab- 
sorbed, and  returned  to  the  chambers  through  the  intervention  of 
the  Glover  tower  G,  as  above  described. 

The  acid  which  collects  in  the  chambers  is  usually  not  permitted 
to  reach  a  higher  specific  gravity  than  about  1.6,  when  it  contains 
about  68  per  cent,  of  sulphuric  acid  ;  for  if  the  strength  be  allowed 
to  exceed  this,  the  acid  not  only  begins  to  dissolve  the  nitrogen 
peroxide  in  the  chamber,  but  exerts  a  corrosive  action  upon  the 
lead  of  which  the  chamber  is  constructed.  It  is  therefore  with- 
drawn, and  the  first  stage  in  the  further  concentration  is  effected 
either  by  the  action  of  the  Glover  tower,  or  by  evaporation  in 
shallow  leaden  pans. 

In  order  to  bring  up  the  strength  of  the  acid  to  that  of  "  oil  of 
vitriol,"  that  is,  to  about  98  per  cent.,  the  acid  from  the  Glover 
tower  or  the  leaden  pans  is  heated  in  either  glass  or  platinum 
stills. 

Sulphuric  acid,  unless  specially  purified,  is  liable  to  contain  a 
number  of  impurities,  such  as  lead  sulphate,  derived  from  the 
action  of  the  acid  upon  the  chamber  ;  arsenic,  from  the  pyrites 
employed  ;  oxides  of  nitrogen,  and  sulphur  dioxide.  From  most 
of  the  impurities,  except  the  arsenic,  the  acid  may  be  purified 
by  the  addition  of  ammonium  sulphate,  and  subsequent  redis- 
tillation— 


Arsenic  is  removed  by  precipitation  of  the  sulphide,  by  means 
of  sulphuretted  hydrogen,  from  the  acid  in  a  moderately  dilute 
state. 


432  Inorganic  Chemistry 

(5.)  The  "  Contact  Process"  As  already  stated,  when  a  mixture 
of  sulphur  dioxide  and  oxygen  is  brought  into  contact  with  finely 
divided  platinum,  the  metal  acts  the  part  of  a  carrier,  or  catalytic 
agent,  and  causes  the  union  of  the  gases.  By  absorbing  the  sulphur 
trioxide  so  produced  in  water,  sulphuric  acid  is  obtained. 

The  chief  obstacles  to  the  successful  utilisation  of  this  re- 
action on  a  manufacturing  scale  are  the  impurities  present  in 
the  sulphur  dioxide  derived  from  the  pyrites  burners,  which  are 
found  to  rapidly  destroy  the  effectiveness  of  the  platinum.  By  the 
system  of  purification  now  adopted  these  difficulties  have  been 
removed,  and  the  operation  is  being  conducted  on  a  successful 
manufacturing  scale. 

In  outline  the  process  is  the  following.  The  mixture  of  sulphur 
dioxide  and  air  drawn  from  the  pyrites  burners  is  first  passed 
through  a  chamber  called  the  "dust  chamber,"  into  which  jets  of 
steam  can  be  injected.  This  serves  the  twofold  purpose  of  remov- 
ing dust  carried  mechanically  from  the  burner,  and  of  diluting  and 
partially  removing  the  sulphuric  acid  which  is  also  a  product  of  the 
burner.  The  gases  after  being  sufficiently  cooled  are  then  made 
to  pass  up  through  a  series  of  towers  (resembling  Glover  towers) 
where  they  meet  a  descending  spray  of  water.  They  are  next 
dried  by  passing  up  another  tower  (which  may  be  compared  to  the 
Gay-Lussac  tower),  where  they  meet  a  descending  stream  of  strong 
sulphuric  acid.  The  gases  are  then  admitted  to  the  contact  chamber, 
which  consists  of  a  vessel  containing  a  number  of  small  perforated 
shelves  upon  which  is  spread  a  layer  of  platinised  asbestos.*  The 
shelves  are  arranged  one  above  the  other  in  tall  narrow  columns 
which  are  separated  from  each  other  in  the  chamber  by  a  slight 
space,  the  object  being  to  prevent  the  mass  from  locally  overheating 
during  the  operation. 

At  the  commencement  the  vessel  is  gently  heated  by  gas  jets, 
but  when  the  operation  has  once  started  external  heat  is  with- 
drawn, and  care  is  then  required  to  prevent  the  temperature  rising 
above  about  350°  (which  is  found  to  be  the  most  favourable  tempera- 
ture) owing  to  the  heat  of  union  of  sulphur  dioxide  and  oxygen. 

Properties. — Sulphuric  acid  is  a  perfectly  colourless,  heavy, 
oily  liquid.  The  acid  obtained  ,by  distillation  always  contains 
about  2  per  cent,  of  water  ;  stronger  than  this  it  cannot  be  prepared 

*  Ferric  oxide  may  be  substituted  for  platinum,  but  the  percentage  yield  is 
smaller. 


Sulphuric  Acid  433 

by  evaporation  or  distillation.  When,  however,  acid  of  this  strength 
is  cooled  to  o°,  colourless  crystals  of  pure  sulphuric  acid,  containing 
100  per  cent.  H2SO4,  are  deposited.  The  crystals  melt  at  10.5°, 
and  remain  liquid  at  temperatures  much  below  this  point.  The 
specific  gravity  of  the  pure  acid  is  1.854  at  o°.  When  boiled  it 
gives  off  sulphur  trioxide  until  the  amount  of  water  in  it  rises 
to  1.5  per  cent.,  when  it  distils  unchanged  at  a  temperature  of 

338°. 

Sulphuric  acid  has  a  powerful  affinity  for  water,  and  absorbs 
moisture  from  the  atmosphere  with  great  readiness.  On  this 
account  it  constitutes  one  of  the  most  valuable  desiccating  agents, 
and  is  constantly  made  use  of  for  depriving  gases,  upon  which  it 
exerts  no  chemical  action,  of  aqueous  vapour.  Owing  to  its  strong 
affinity  for  water  it  decomposes  many  organic  substances  contain- 
ing hydrogen  and  oxygen,  withdrawing  from  the  compounds  these 
elements  in  the  proportion  to  yield  water  ;  its  action  upon  formic 
acid,  oxalic  acid  (see  Carbon  Monoxide),  and  alcohol  (see  Ethy- 
lene)  are  examples  of  this  action. 

When  the  acid  is  poured  upon  such  substances  as  wood  or  sugar 
the  elements  composing  water  are  withdrawn  and  the  carbon  is 
liberated,  with  the  result  that  the  compounds  are  blackened  or 
charred. 

When  sulphuric  acid  is  mixed  with  water  considerable  heat  is 
disengaged,  the  temperature  often  rising  to  the  boiling-point  of 
water,  and  at  the  same  time  a  diminution  in  volume  takes  place. 
The  maximum  contraction  is  obtained  upon  mixing  the  materials 
in  the  proportion  of  one  molecule  of  acid  to  two  molecules  of  water. 
The  diminution  in  volume  in  this  case  amounts  to  8  per  cent.,  and 
the  composition  of  the  acid,  produced  corresponds  to  the  formula 
H2SO4,2H2O. 

Sulphuric  acid  combines  with  water  in  various  proportions,  form- 
ing a  number  of  hydrates  of  a  more  or  less  definite  character.  The 
best  known  are  those  represented  by  the  formulae  H2SO4,H2O  and 
H2SO4,2H2O.  These  compounds  may  be  regarded  as  respectively 
tetrabasic  and  hexabasic  sulphuric  acid,  and  their  relation  to 
the  ordinary  dibasic  acid  may  be  expressed  by  the  following 
formulae  — 


2  E 


HoSO4        .        .        .     orS 
H4S05  or  H9S04,H20     „  SO(HO)4. 
H6S06  „  H2S04,2H20   „  S(HO), 


4.34  Inorganic  Chemistry 

Salts  of  each  of  these  acids  are  known — 

Hydrogen  potassium  sulphate  .     .  HKSOi  \ 

Normal  potassium  sulphate      .     .  K2SO4    >  Derived  from  H2SO4. 

Barium  sulphate BaSO4     ) 

Tetrabasic  lead  sulphate       .     .     .  Pb2SO5  „          „     H4SO5. 

Hexabasic  mercuric  sulphate  )  TT_  Qn  u  Qr» 

(Turpeth  mineral)  }      '  Hg3S°6  "          »     H&°* 

Most  sulphates  are  soluble  in  water  ;  those  of  lead,  calcium,  and 
strontium  are  only  very  sparingly  soluble,  whilst  barium  sulphate  is 
insoluble  both  in  water  and  acids.  The  presence  of  sulphuric  acid 
or  a  sulphate  may  therefore  be  readily  detected  by  the  addition  of 
a  soluble  barium  salt,  which  causes  the  immediate  precipitation  of 
white  barium  sulphate,  insoluble  in  hydrochloric  acid. 


PYROSULPHUEIC  ACID  (Nordhausen  Acid  ;  Fuming  Sulphuric  Acid]. 


Formula,  H2S2O7  or     oo 

Modes  of  Formation.—  (  i.)  This  acid  may  be  obtained  by  dis- 
solving sulphur  trioxide  in  ordinary  sulphuric  acid  — 

H2SO4  +  S03  =  H2S2O7. 

On  cooling  the  solution  to  o°  the  pyrosulphuric  acid  separates  out 
in  the  form  of  large  colourless  crystals. 

(2.)  Pyrosulphuric  acid  is  manufactured  by  the  distillation  of 
ferrous  sulphate  in  clay  retorts,  mounted  in  series  in  a  large 
"  galley  "  furnace.  The  first  action  of  heat  upon  crystallised  ferrous 
sulphate  (green  vitriol)  is  to  expel  six  molecules  of  water  of  crystal- 
lisation, leaving  the  salt  of  the  composition  FeSO4,H2O.  When 
this  substance  is  further  heated  it  is  decomposed  finally  into  ferric 
oxide,  with  the  formation  of  sulphur  trioxide,  water,  and  sulphur 
dioxide,  thus  — 

2FeSO4,H2O  =  Fe2O3  +  SO3  +  SO2  +  2H2O. 

The  decomposition  takes  place  in  two  stages,  the  sulphur  dioxide 
and  water  being  evolved  in  the  first  part  of  the  process  with  the 
formation  of  ferric  sulphate,  which  is  afterwards  broken  up  in  the 
manner  shown  in  the  following  equation  — 


(i.) 

(2.)  Fe2(S04)3=Fe20 


Thiosulphuric  Acid  435 

The  sulphur  trioxide  is  condensed  in  receivers  containing  either 
a  small  quantity  of  water  or  a  charge  of  sulphuric  acid. 

(3.)  Pyrosulphuric  acid  may  also  be  obtained  by  decomposing 
sodium  pyrosulphate  (Na2S2O7),  either  by  heating  it  to  a  high 
temperature  (see  Sulphur  Trioxide,  page  422),  or  by  acting  upon  it 
with  sulphuric  acid,  thus  — 

Na2S2O7  +  H2SO4=2HNaSO4+SO3. 

The  sulphur  trioxide  obtained  is  dissolved  in  sulphuric  acid,  as  in 
the  former  methods  ;  and  the  hydrogen  sodium  sulphate,  when 
gently  heated  to  about  300°,  is  reconverted  into  pyrosulphate  by 
the  loss  of  a  molecule  of  water  (page  422). 

Properties.—  Pyrosulphuric  acid  is  a  colourless,  strongly  fuming 
liquid,  having  a  specific  gravity  of  1.88.  When  cooled,  it  solidifies 
to  a  crystalline  mass,  which  melts  at  35°.  The  compound  may  be 
regarded  as  consisting  of  one  molecule  of  sulphuric  acid  plus  a 
molecule  of  sulphur  trioxide,  H2SO4,SO3  ;  or,  as  being  derived 
from  two  molecules  of  sulphuric  acid,  by  the  withdrawal  of  one 
molecule  of  water,  thus  — 


Pyrosulphuric  acid  forms  a  stable  series  of  salts,  of  which  the 
sodium  compound  already  mentioned  is  a  typical  example.  These 
salts  are  sometimes  spoken  of  as  the  disulphates,  and  are  analogous 
to  the  dichromates  (q-'V^ 

Two  other  definite  compounds  of  sulphur  trioxide  and  sulphuric  acid  are 
known  to  exist,  both  of  which  are  fuming  acids.  The  composition  of  these 
substances  is  expressed  by  the  formulae— 

H2S04,3S03,  or  H2S4O13  ;  and  3H2SO4,SO3,  or  HfiS4O15. 


THIOSULPHURIC  ACID. 

Formula,  H2S2O3. 

This  acid  has  never  been  obtained  in  the  free  state,  as  it  decom- 
poses almost  as  soon  as  liberated  from  its  salts  into  sulphur  dioxide 
and  water,  with  precipitation  of  sulphur — 

H2S2O3  =  SO2  +  H2O  +  S. 


436  Inorganic  Chemistry 

The  thiosulphates,  however,  are  stable  and  important  salts,  the 
sodium  salt  being  largely  used  in  photography  under  the  name  of 
hyposulphite  of  soda,  or  "  hypo." 

Modes  cf  Formation  of  Thiosulphates.—  (i.)  These  salts 
may  be  obtained  by  digesting  flowers  of  sulphur  with  solutions  of 
the  sulphites,  thus  — 

S  =  Na2S2O3. 

(2.)  Sodium  thiosulphate  is  also  formed  when  sulphur  dioxide  is 
passed  into  a  solution  of  sodium  sulphide.  The  reaction  may  be 
regarded  as  taking  place  in  three  steps,  in  which  sodium  sulphite 
and  sulphuretted  hydrogen  are  the  first  products.  The  latter  com- 
pound is  then  acted  upon  by  sulphur  dioxide,  with  the  precipitation 
of  sulphur,  thus  — 

SO2  +  H2O  +  Na2S  =  Na2SO3  +  H2S. 


And  the  sulphur  reacts  with  the  already  formed  sulphite,  as  indi- 
cated in  the  equation  given  above. 

(3.)  When  sulphur  is  boiled  with  sodium  hydroxide,  or  with  milk 
of  lime,  mixtures  of  sulphides  and  thiosulphates  are  obtained  in 
both  cases— 


3Ca(HO)2  +  12S  =  CaS2O3  +  2CaS6  +  3H2O. 

The  sodium  sulphide  can  be  converted  into  thiosulphate  by  the 
reactions  given  above.  Calcium  pentasulphide,  on  exposure  to  air, 
absorbs  oxygen  and  forms  a  further  quantity  of  thiosulphate  with 
precipitation  of  sulphur  — 

CaS5  +  3O  =  3S  +  CaS2O3. 

The  thiosulphates  are  decomposed  by  most  acids,  with  the  libera- 
tion of  sulphur  dioxide,  and  precipitation  of  sulphur.  They  show  a 
great  tendency  to  form  double  salts,  many  of  which  are  soluble 
in  water  ;  thus  sodium  thiosulphate,  in  contact  with  either  silver 
chloride,  bromide,  or  iodide,  forms  the  soluble  double  sodium-silver 
thiosulphate,  NaAgS2O3— 

Na2S2O3  +  AgCl  =  NaCl  +  NaAgS2O3. 


Trithionic  Acid  437 

The  employment  of  sodium  thiosulphate  in  photography,  for 
"fixing"  negatives  or  silver  prints,  depends  upon  this  property. 

Thiosulphuric  acid  may  be  regarded  as  being  derived  from  sul- 
phuric acid  by  the  replacement  of  one  of  the  (HO)  or  hydroxyl 
groups,  by  an  equivalent  of  (HS)  or  hydrosulphyl  — 

H0?so  HS 

no;50*-     Ho 

Dithionic  Acid,  H2S2O6  or  Holo2  }'~This  comPound  is  prepared  by 
passing  a  stream  of  sulphur  dioxide  through  water  in  which  manganese 
dioxide  is  suspended,  whereby  manganese  dithionate  is  formed  ;  while  at  the 
same  time  a  portion  of  the  salt  is  acted  upon  by  manganese  dioxide,  and  con- 
verted into  manganous  sulphate,  thus  — 

2SO2+MnO2=MnS2O6. 
=2MnSO4. 


On  the  addition  of  barium  hydroxide  to  the  solution,  barium  dithionate, 
barium  sulphate,  and  manganous  hydrate  are  formed  — 

MnS206+Ba(HO)2=BaS206+Mn(HO)2. 

Barium  dithionate,  being  soluble,  is  separated  by  nitration,  and  upon 
evaporation  separates  out  in  crystals  of  the  composition  BaS2O6,2H2O. 

Upon  the  addition  of  dilute  sulphuric  acid  in  amount  demanded  by  the 
equation  — 

BaS206  +  H2S04  =  BaS04  +  H2S2O6, 

the  acid  itself  is  obtained.  The  solution  may  be  concentrated  in  vacuo  until 
it  reaches  a  specific  gravity  of  1.347.  Further  concentration  results  in  its  de- 
composition into  sulphuric  acid  and  sulphur  dioxide  — 

H2S206=S02+H2S04. 

Dithionic  acid  forms  well-defined  crystalline  salts,  which  on  heating  decom- 
pose into  sulphates  with  evolution  of  sulphur  dioxide. 

Dithionic  acid  was  formerly  called  hyposulphuric  acid,  and  its  salts  are  still 
sometimes  referred  to  as  hyposulphates. 

TritMonic  Acid,  H2S3O6  or  HO'SQ2  }S'~The  Potassium  salt  of  this  acid 
may  be  obtained  by  passing  sulphur  dioxide  through  a  strong  solution  of 
potassium  thiosulphate— 

3S02  4-  2K2S2O3  =  S  +  2K2S3O6. 
It  is  also  formed  when  a  solution  of  potassium  silver  thiosulphate  is  boiled— 

KO 


438  Inorganic  Chemistry 

The  sodium  salt  may  be  obtained  by  the  addition  of  iodine  to  a  mixture  of 
sodium  sulphite  and  thiosulphate — 

NaO  ) 

NaS  I 

The  acid  itself  is  obtained  by  the  addition  of  fluosilicic  acid  to  a  solution  of 
the  potassium  salt,  when  insoluble  potassium  fluosilicate  is  precipitated. 

Both  the  acid  itself  and  its  salts  are  readily  decomposed  into  sulphur  dioxide, 
sulphur,  and  either  sulphuric  acid  or  a  sulphate,  thus — 

K2S3O6=K2SO4+SO2+S. 

When  acted  upon  by  sodium  amalgam,  sodium  trithionate  is  converted 
back  again  into  its  generators,  sodium  sulphite  and  thiosulphate,  thus— 

NaO'SOo)^  NaO  )  „, 


TetratMonic  Acid,  H2S4O6or         -2    s2.—  The  sodium  salt  is  obtained 


by  the  action  of  iodine  upon  sodium  thiosulphate  — 


The  barium  salt,  from  which  the  acid  itself  is  most  readily  obtained,  is  pre- 
pared by  the  gradual  addition  of  iodine  to  barium  thiosulphate  in  water  — 

2BaS2O3+I2=BaI2  +  BaS4O6. 

The  barium  tetratbionate  is  separated  by  the  addition  of  alcohol,  which  dis- 
solves the  iodide  and  excess  of  iodine,  leaving  the  tetrathionate.  By  the 
addition  of  dilute  sulphuric  acid  to  an  aqueous  solution  of  this  salt,  in 
amount  demanded  by  the  equa'tion  — 

BaS406  +  H2SO4  =  H2S406  +  BaSO4, 

a  dilute  aqueous  solution  of  the  acid  may  be  obtained.  The  dilute  acid  may 
be  boiled  without  decomposition  ;  but  when  concentrated,  it  readily  passes 
into  sulphuric  acid,  sulphur  dioxide,  and  sulphur. 

Sodium  amalgam  decomposes  the  sodium  salt  into  two  molecules  of  thio- 
sulphate, reversing  the  reaction  by  which  it  is  produced. 


Pentathionic  Acid,   H2S5O6  or  2    s3.—  This  acid   is  prepared  by 

"          * 


passing  sulphuretted  hydrogen   into  a   strong  aqueous   solution  of  sulphur 
dioxide  — 

5SO2+5H2S=HoS5Oe+5S  +  4HoO. 
or  — 

5H2SO:{+5H2S=H2S506  +  5S  +  9H20. 

The  solution  contains,  however,  more  or  less  of  the  other  thionic  acids,  but 
«s  the  passage  of  sulphuretted  hydrogen  is  continued,   these   are  gradually 


Oxychlorides  of  Sulphur  439 

decomposed,  and  ultimately  the  pentathionic  acid   also,   so  that   the   final 
products  of  the  action  of  excess  of  this  gas  will  be  sulphur  and  water — 

H2S5O6  +  5H2S=6H2O  +  10S. 

The  solution  obtained  by  the  first  action  may  be  concentrated  by  cautious 
evaporation  in  vacuo,  until  a  specific  gravity  of  1.46  is  obtained,  when  on  partial 
saturation  with  potassium  hydroxide  and  filtration,  a  solution  is  obtained  which 
on  spontaneous  evaporation  deposits  crystals  of  potassium  pentathionate,  having 
the  composition  K2S5O6,3H2O.  On  heating,  the  salt  splits  up  into  potassium 
sulphate,  sulphur  dioxide,  and  sulphur. 


OXYCHLORIDES   OF   SULPHUR. 

Four  of  these  compounds  are  known,  all  of  which  may  be  re- 
garded as  being  derived  from  the  oxyacids  by  the  replacement  of 
hydroxyl  (HO)  by  its  equivalent  of  chlorine. 

1.  Thionyl  chloride,  or   Cl )  ,,~  . ,  HO  )  ,  „  sulphurous 
c   /^t         •    L7     -j     /-,  t  SO  corresponding  oxyacidTT~  I  SO, 
Sulphurous  chloride    Cl)                                       J         HO  j  acid. 

2.  Sul phuryl  chloride,  or  Cl  )  CQ 

Sulphuric  chloride     Cl  j       2  )  HO  )  ~n    sulphuric 

3.  Sulphuric  chlorhydrate,  or  Cl)  SQ  j  HO  I       2>      acid. 
Chlorosulphonic  acid        HO  J 

4.  Disulphuryl  chloride,  or  C1'SO2  \~  HO'SO2  )  Q  pyrosulphu- 
Pyrosulphuric  chloride  C1'SO2  j  "                "HO'SO2  I     '     ric  acid. 

Thionyl  Chloride,  SOCLj,  molecular  weight =118.96,  is  obtained  by  the 
action  of  phosphorus  pentachloride  upon  sodium  sulphite— 

SO(NaO)2+2PCl5=SOCl2  +  2POCl3  +  2NaCl. 

It  is  also  obtained  when  dry  sulphur  dioxide  is  passed  over  phosphorus 
pentachloride — 

;so2+pci5=soci2+poci3. 

Properties. — Thionyl  chloride  is  a  colourless  and  highly  refractive  liquid 
which  fumes  in  moist  air,  and  has  a  pungent  unpleasant  smell.  It  boils  at  78°, 
and  is  at  once  decomposed  by  water  into  its  corresponding  oxyacid,  with  for- 
mation of  hydrochloric  acid — 

SOC12+2H2O=H2SO3+HC1. 

Sulphuryl  Chloride,  SO2C12;  molecular  weight  =  134. 96.  This  compound 
(sometimes  known  as  chlorosulphuric  acid]  can  be  obtained  by  the  direct 
union  of  chlorine  and  sulphur  dioxide,  under  the  prolonged  influence  of  bright 
sunlight — 

S02+C12=S02C12. 

It   is  also  formed  by  the  action  of  heat  upon    sulphuric  chlorhydrate. 


44O  Inorganic  Chemistry 

This  substance,  on  being  simply  heated  to  180°  in  sealed  tubes  for  a  few  hours, 
breaks  up  into  sulphuryl  chloride  and  sulphuric  acid  — 


0  <5n  Qn  0 

2  C1    jSO2=ajS02+HOjS02. 

Properties.  —  Sulphuryl  chloride  is  a  colourless  liquid,  which  fumes  in  moist 
air,  and  has  a  specific  gravity  of  1.66.  It  boils  at  70°,  and  is  decomposed  by 
water,  with  formation  of  sulphuric  acid  and  hydrochloric  acid  — 


Q  }-S02  +  2H20=2HCl  +  ^8  }s°2- 


Sulphuric  Chlorhydrate,  SO2C1(HO).—  This  compound  is  the  first  pro- 
duct of  the  replacement  of  the  (HO)  groups  in  sulphuric  acid  by  chlorine, 
and  is  formed  by  the  direct  combination  of  sulphur  trioxide  and  hydrochloric 
acid  — 

SO3+HC1=HC1SO3  or  SO2C1(HO). 

It  may  be  obtained  by  distilling  sulphuric  acid  with  phosphorus  oxychloride— 


Or  by  passing  dry  gaseous  hydrochloric  acid  into  melted  pyrosulphuric  acid  — 
H2S2O7  +  2HC1=H2O  +  2HC1SO3. 

Properties.  —  Sulphuric  chlorhydrate  is  a  colourless  fuming  liquid,  having  a 
specific  gravity  of  1.76,  and  boiling  at  i49°-i5i°,  with  partial  dissociation  into 
its  generators,  sulphur  trioxide  and  hydrochloric  acid.  In  contact  with  water 
it  is  decomposed  with  considerable  violence,  with  formation  of  sulphuric  and 
hydrochloric  acids  — 


Disulphuryl  Chloride  (pyrosulphuric  chloride},  QSOV  °  or  S2°5C^-  This 
substance  is  obtained  by  the  action  of  sulphur  trioxide,  or  sulphuric  chlor- 
hydrate, upon  phosphorus  pentachloride  — 

2SO3  +  PC15  =  POC13  +  S205C12. 
2S02C1(  HO)  +  PC15=  POC13  +  2HC1  +  S2O5C12. 

It  is  also  produced  by  the  action  of  sulphur  trioxide  upon  sulphur 
dichloride  — 

5SO3  +  S2C12  =  S205C12  +  5S02. 

Or  by  the  action  of  sulphur  trioxide  upon  sulphuric  chloride  — 


Carbon  Bisulphide  441 

Properties. — Pyrosulphuric  chloride  is  a  heavy,  oily,  fuming  liquid,  resem- 
bling pyrosulphuric  acid  in  appearance.  It  has  a  specific  gravity  of  1.819,  anc* 
boils  at  146°.  When  mixed  with  water  it  slowly  decomposes  into  sulphuric 
and  hydrochloric  acids,  showing  a  marked  difference  in  this  respect  from 
sulphuric  chlorhydrate — 

S205C12+3H20=2H2S04+2HC1. 

COMPOUNDS  OF  SULPHUR  WITH  FLUORINE. 

Perfluoride  of  Sulphur,  SF6.— This  compound  has  been  recently  obtained 
(Moissan)  by  passing  fluorine  over  sulphur. 

Properties. — Sulphur  perfluoride  is  a  colourless  inodorous  gas,  very  soluble 
in  water,  and  incombustible  in  air.  It  is  a  comparatively  inactive  compound. 

Thionyl  Fluoride,  SOF2,  is  obtained  by  the  action  of  fluorine  upon  thionyl 
chloride  ;  also  by  the  action  of  arsenic  trifluoride  upon  thionyl  chloride — 

2AsF3+3SOCl2=3SOF2  +  AsCl3. 

Properties. — Thionyl  fluoride  is  a  colourless  gas  which  fumes  strongly  in 
moist  air.  It  is  immediately  decomposed  by  water— 

SOF2+H2O=2HF+SO2. 


CARBON  BISULPHIDE. 

Formula,  C S2.     Molecular  weight =76. 12.     Vapour  density =38. 06. 

History. — This  compound  was  accidentally  produced  by  Lam- 
padius  (1796)  when  heating  a  mixture  of  charcoal  and  pyrites. 

Mode  Of  Formation. — Carbon  disulphide  is  prepared  by  passing 
the  vapour  of  sulphur  over  red-hot  charcoal,  when  the  two  elements 
unite  and  form  the  volatile  product,  which  is  condensed  in  vessels 
surrounded  with  cold  water — 

C  +  S2  =  CS2. 

The  product  is  always  contaminated  with  free  sulphur,  which 
volatilises,  and  is  also  accompanied  by  considerable  quantities  of 
sulphuretted  hydrogen,  formed  by  the  action  of  sulphur  upon  the 
hydrogen  contained  in  the  charcoal. 

When  carbon  disulphide  is  prepared  on  a  manufacturing  scale, 
the  charcoal  is  heated  in  a  vertical  cast-iron  or  earthenware  retort, 
C  (Fig.  117),  having  an  elliptical  section,  and  provided  with  three 
openings.  The  retort  is  built  into  a  suitable  furnace,  whereby  it 
can  be  uniformly  heated  to  redness.  A  quantity  of  sulphur,  con- 
tained in  the  pot  S,  kept  liquid  by  the  heat  of  the  furnace,  is 


442 


Inorganic  Chemistry 


allowed  to  enter  at  intervals  through  the  pipe  B.  As  the  vapour 
comes  in  contact  with  the  red-hot  charcoal,  combination  ensues, 
and  the  carbon  disulphide  escapes  through  the  pipe  D,  which  is 
inclined  to  the  retort  so  as  to  allow  condensed  sulphur  to  run  back. 
Sulphur  which  escapes  condensation  in  this  pipe,  collects,  for  the 
most  part,  in  the  vessel  E,  which  is  closed  by  water  seals,  as  seen 
in  the  figure.  The  volatile  compounds  are  then  passed  through  a 
Liebig's  condenser  about  30  ft.  long,  and  the  crude  disulphide  so 
condensed  is  collected  in  a  receiver.  Any  vapour  of  carbon  disulphide 

which  is  carried  on  by  the  sul- 
phuretted hydrogen  is  absorbed 
by  passing  the  gas  through  a 
scrubber  containing  oil  ;  and 
finally  the  sulphuretted  hydrogen 
is  absorbed  in  a  lime  purifier, 
similar  to  those  employed  for  the 
purification  of  coal  gas.  The 
ashes  are  withdrawn  from  the 
retort  through  the  wide  tube  B  ; 
and  the  fresh  charcoal  is  intro- 
duced through  the  opening  A. 
In  order  to  prevent  the  escape 
of  the  unpleasant  and  injurious 
vapours  from  A  during  the  addi- 
tion of  fresh  charcoal,  the  opening 
A'  is  put  into  communication  with  the  chimney  of  the  furnace. 
The  sulphur  which  flows  back  into  the  retort  from  D  is  conveyed, 
by  means  of  the  pipe^j  nearly  to  the  bottom  of  the  mass  of  heated 
charcoal,  so  that  its  vapour  shall  once  more  be  made  to  pass  over 
the  carbon. 

At  the  present  day,  since  the  application  of  electrical  heating  to 
manufacturing  processes,  the  mixture,  instead  of  being  heated  from 
the  outside  by  fuel,  is  heated  inside  in  vessels  of  modified  form  by 
means  of  the  electric  arc. 

The  crude  product  is  purified  by  distillation  and  subsequent 
agitation  with  mercury. 

Properties. — Carbon  disulphide  is  a  colourless,  mobile,  and 
highly  refracting  liquid.  When  perfectly  pure  it  possesses  a 
sweetish,  and  not  unpleasant,  ethereal  smell,  but  as  usually  met 
with  the  odour  is  decidedly  foetid. 

Its  specific  gravity  at  o°  is  1.292,  and  it  boils  at  46°.     The  vapour 


FIG.  117. 


Carbon  Bisulphide  443 

of  carbon  disulphide  has  a  very  low  igniting-point  (see  page  329). 
It  burns  with  a  blue  flame,  which,  when  fed  with  oxygen,  emits  a 
dazzling  blue  light.  When  carbon  disulphide  vapour  is  mixed 
with  three  times  its  volume  of  oxygen,  and  a  light  applied,  the 
mixture  explodes  with  violence  ;  the  products  of  the  combustion 
being  carbon  dioxide  and  sulphur  dioxide  — 


The  vapour  of  carbon  disulphide,  when  constantly  inhaled  in 
small  quantities,  has  an  injurious  effect  upon  the  health,  and  if 
breathed  in  large  quantities  is  a  powerful  poison. 

When  heated  to  a  bright  red  heat,  carbon  disulphide  vapour  is 
decomposed  into  its  constituent  elements  :  on  this  account,  in  the 
manufacture  of  this  compound,  care  is  taken  that  the  temperature 
does  not  rise  too  high. 

The  vapour  of  carbon  disulphide  is  decomposed  by  potassium, 
which,  when  heated,  burns  in  the  vapour,  forming  potassium 
sulphide,  and  liberating  carbon  — 

CS2  +  2K2  =  C  +  2K,S. 

When  passed  over  heated  slaked  lime,  carbon  disulphide  vapour 
is  converted  into  carbon  dioxide  and  sulphuretted  hydrogen  — 


This  reaction  is  made  use  of  for  converting  the  carbon  disul- 
phide, which  is  always  present  in  coal  gas,  into  the  two  easily 
removed  substances,  carbon  dioxide  and  sulphuretted  hydrogen. 

When  a  mixture  of  carbon  disulphide  vapour  and  sulphuretted 
hydrogen  is  passed  over  heated  copper,  marsh  gas  is  formed  — 


Carbon  disulphide  is  soluble  to  a  minute  extent  in  water  ;  I 
volume  of  water  dissolves  .001  volume  of  this  liquid,  and  the 
solution  possesses  the  taste  and  the  smell  of  the  disulphide.  It 
mixes  in  all  proportions  with  alcohol,  ether,  the  hydrocarbons  of 
the  benzene  family,  and  most  essential  oils.  It  also  dissolves 
sulphur,  phosphorus,  iodine,  bromine,  caoutchouc,  and  most  fats  ; 
and  is  largely  used  in  the  arts,  both  as  a.  solvent  for  caoutchouc 
and  in  extracting  essential  oils,  spices,  and  perfumes. 

Thioearbonie  Acid.—  Carbon  disulphide  is  the  sulphur  ana- 
logue of  carbon  dioxide,  CS2  ;  CO2.  Like  the  oxygen  compound, 


444  Inorganic  Chemistry 

it  forms  a  feeble  acid,  which  has  received  the  name  thiocarbonic 
acid,  H2CS3  ;  carbonic  acid,  H2CO3. 

The  thiocarbonates  are  produced  by  reactions  analogous  to 
those  by  which  carbonates  are  formed.  Thus,  when  carbon  disul- 
phide  is  brought  into  contact  with  potassium  sulphide,  potassium 
thiocarbonate  is  obtained  — 

C  82  -f"  K2S  =  KoC  S3. 

Thiocarbonates  are  likewise  formed  by  the  action  of  carbon 
disulphide  upon  metallic  hydroxides  — 


The  acid  itself  is  obtained  as  a  yellow  oil,  having  an  unpleasant 
odour  by  the  decomposition  of  a  thiocarbonate  by  dilute  hydro- 
chloric acid. 

A  large  number  of  compounds  are  known  in  which  divalent 
sulphur  replaces  oxygen,  and  which  therefore  stand  in  the  same 
relation  to  the  oxygen  compounds  as  thiocarbonic  acid  stands  to 
carbonic  acid  ;  for  example  — 

Thiocarbamic  acid,  CS2,NH3,  or  -rrq 


Carbamic  acid,  CO2,NH3,  or          2  \  CO. 

Other  Compounds  of  Carbon  and  Sulphur.—  When  carbon  disulphide  is 
exposed  to  the  influence  of  light,  there  is  gradually  formed  upon  the  glass 
vessel  containing  it  a  brown  deposit,  which  is  believed  to  be  carbon  mono- 
sulphide,  CS  ;  the  sulphur  analogue  of  carbon  monoxide.  When  electric 
sparks  from  carbon  poles  are  passed  through  the  vapour  of  carbon  disulphide, 
or  when  the  electric  arc  is  produced  in  the  vapour,  an  offensive  smelling  liquid 
is  obtained,  which  exerts  a  most  irritating  and  tear-producing  effect  upon  the 
eyes.  This  liquid  has  been  shown  to  have  the  composition  C3S2.* 


SELENIUM. 

Symbol,  Se.     Atomic  weight =79.1.     Molecular  weight  =  158. 2. 

History. — This  element  was  discovered  by  Berzelius  (1817),  who  gave  it 
the  name  selenium  (signifying  the  moon)  on  account  of  its  close  analogy  with 
the  previously  discovered  element  tellurium  (signifying  the  earth). 

Occurrence.— Selenium  is  occasionally  met  with  associated  with  native 
sulphur,  probably  as  a  selenide  of  sulphur.  In  a  few  minerals  of  considerable 

*  Von  Lengyel,  1894. 


Selenium  445 

rarity,  selenium  is  met  with  in  the  form  of  selenides  of  such  metals  as  mercury, 
lead,  silver.  It  occurs  in  very  small  quantities  in  a  large  number  of  metallic 
sulphides. 

Modes  of  Formation.— ( i.)  When  pyrites  containing  selenium  is  employed 
in  the  manufacture  of  sulphuric  acid,  the  selenium  is  oxidised  by  the  atmos- 
pheric oxygen  into  selenium  dioxide,  which  is  carried  forward  with  the  sulphur 
dioxide.  Selenium  dioxide,  being  a  solid  substance,  is  partly  deposited  in  the 
flues,  and  in  the  Glover  tower,  and  partly  carried  forward  into  the  chambers, 
where  it  forms  a  red-coloured  deposit.  To  obtain  the  selenium,  either  the  flue 
dust  or  the  chamber  deposit  is  first  boiled  with  dilute  sulphuric  acid,  and  either 
nitric  acid  or  potassium  chlorate  added,  in  order  to  oxidise  it  completely  into 
selenic  acid,  H2SeO4.  The  solution  is  then  boiled  with  strong  hydrochloric 
acid,  whereby  it  is  reduced  to  selenious  acid,  H2SeO3,  when  a  stream  of  sulphur 
dioxide  is  passed  through  it  which  precipitates  the  selenium  as  a  red  powder — 

H2SeO3  +  2SO2  +  H2O = Se  +  2H2SO4. 

(2.)  A  second  method  for  the  preparation  of  selenium  from  the  chamber 
deposit  consists  in  digesting  the  substance  with  potassium  cyanide,  whereby 
it  is  converted  into  soluble  potassium  selenocyanide,  SeK(CN).  On  the 
addition  of  hydrochloric  acid  to  this  solution,  the  element  is  precipitated  as  a 
red  amorphous  powder,  and  hydrocyanic  acid  and  potassium  chloride  go  into 
solution — 

SeK(CN)  +  HCl=Se  +  KC1  +  H(CN). 

Properties. — Selenium  is  known  in  various  allotropic  modifications. 

1.  Soluble  in  carbon  disulphide. — o.  Brick-red  amorphous  powder,  obtained 
by  precipitation  with  acids,  or  reduction  of  selenious  acid,  in  the  cold,  by 
sulphur  dioxide. 

/3.  Black  crystalline  powder,  obtained  by  reduction  of  hot  selenious  acid  by 
sulphur  dioxide. 

7.  Dark  red  translucent  monoclinic  crystals,  specific  gravity  4.5,  deposited 
from  solution  in  carbon  disulphide. 

5.  Black,  shining,  brittle  amorphous  mass,  having  a  conchoidal  fracture, 
and  a  specific  gravity  of  4.3,  obtained  by  rapidly  cooling  melted  selenium. 

2.  Insoluble  in  carbon  disulphide. — Black,  metallic-looking  crystalline  mass, 
having  a  granular  fracture.     Obtained  by  quickly  cooling  melted  selenium  to 
210°  and  keeping  it  for  some  time  at  that  temperature,  when  the  mass  solidifies 
with  rise  of  temperature  to  217°.     This  insoluble  variety,  sometimes  called 
metallic  selenium,  is  also  formed  as  a  deposit  of  minute  black  crystals,  when 
concentrated  solutions  of  sodium  or  potassium  selenide  are  exposed  to  the  air. 

This  modification  has  a  specific  gravity  of  4.5,  and  melts  at  217°. 

Selenium  boils  at  680°,  forming  a  dark-red  vapour  which  condenses  in  the 
form  of  flowers  of 'selenium ,  having  a  scarlet-red  colour. 

At  high  temperatures  the  vapour  of  selenium,  like  that  of  sulphur,  becomes 
a  true  gas  ;  thus  at  1420°,  the  vapour-density  is  found  to  be  81.5,  approaching 
very  closely  to  the  normal  density  demanded  by  the  molecule  Se2. 

"Metallic"  selenium  conducts  electricity,  and  the  element  exhibits  the 
remarkable  property  of  having  its  conductivity  increased  by  light  ;  the  con- 
ductivity of  selenium  when  exposed  to  diffused  daylight  being  about  twice  as 


446  Inorganic  Chemistry 

great  as  when  in  the  dark.  This  alteration  in  the  electrical  resistance  with 
varying  intensities  of  light,  is  a  property  of  selenium  that  was  made  use  of  in 
the  construction  of  an  instrument  known  as  the  photophone,  but  it  has  not  as 
yet  been  put  to  any  practical  use.  When  selenium  is  heated  in  the  air,  it 
burns  with  a  blue  flame,  with  the  formation  of  selenium  dioxide,  and  at  the  same 
time  emits  a  powerful  and  characteristic  smell  resembling  rotten  horse-radish. 

When  selenium  is  heated  in  a  tube  filled  with  an  indifferent  gas,  it  sublimes 
in  the  form  of  a  red  deposit ;  but  when  heated  in  hydrogen,  the  sublimate  is  in 
the  form  of  black  shining  crystals.  The  formation  of  these  crystals  is  due  to 
the  fact  that  selenium  combines  with  the  hydrogen,  and  the  hydrogen  selenide 
is  again  decomposed  by  the  heat. 

Hydrogen  Selenide  \selenurettedhydrogen),  H2Se  ;  molecular  weight=8i.  12. 
Hydrogen  selenide  is  formed  when  selenium  is  heated  in  hydrogen. 

This  compound  is  also  obtained  by  the  action  of  dilute  hydrochloric  or  sul- 
phuric acid  upon  either  potassium  selenide  or  ferrous  selenide — 

FeSe  +  H2SO4=FeSO4+  H2Se. 

Properties. — Hydrogen  selenide  is  a  colourless  gas,  strongly  resembling 
sulphuretted  hydrogen,  both  in  its  smell  and  in  its  chemical  behaviour.     It  is 
readily  soluble  in  water,  and  when  passed  through  metallic  solutions  precip 
tates  insoluble  selenides  of  most  of  the  heavy  metals.      Hydrogen  selenid 
burns  with  a  blue  flame,  with  the  production  of  water  and  selenium  dioxide. 
Its  smell,  although  resembling  that  of  its  sulphur  analogue,  is  more  unpleasan 
and  its  effects  upon  the  system  are  more  persistent  and  injurious.    A  single  sm; 
bubble  inhaled  through  the  nostril  produces  temporary  paralysis  of  the  olfac- 
tory nerves,  accompanied  by  inflammation  of  the  mucous  membrane. 

No  compound  of  selenium  corresponding  to  hydrogen  disulphide  is  known. 

COMPOUNDS  WITH  HALOGENS. 

Diselenium  Bichloride,  Se2Cl2,  is  obtained  by  passing  chlorine  over 
selenium,  or  by  passing  gaseous  hydrochloric  acid  through  a  solution  of 
selenium  in  nitric  acid. 

Properties. — Selenium  chloride  is  a  brown  oily  liquid,  in  which  selenium 
itself  is  readily  soluble,  and  from  which  the  element  is  deposited  in  the  form 
which  is  insoluble  in  carbon  disulphide.  It  is  slowly  decomposed  by  water, 
thus— 

2Se2Cl2 + 3H20  =  H2SeO3  +  3Se  +  4HC1. 

Corresponding  bromine  and  iodine  compounds  are  known,  Se2Br2,  and 
Se2I2. 

Selenium  Tetrachloride,  SeCl4,  is  prepared  either  by  the  action  of  chlorine 
upon  selenium  chloride — 

Se2Cl2+3Cl2=2SeCl4> 

or  by  heating  a  mixture  of  selenium  dioxide  and  phosphorus  pentachloride — 
3SeO2+3PCl5=3SeCl4-f  P2O5  +  POC13. 

Properties. — Selenium  tetrachloride  is  a  white,  crystalline,  volatile  com- 
pound ;  which  may  be  sublimed  without  decomposition  and  without  fusion. 


Selenic  Acid  447 

When  the  vapour  is  heated  above  200°  it  begins  to  dissociate  into  selenium 
and  chlorine.  It  dissolves  in  water,  with  decomposition  into  hydrochloric  and 
selenious  acids — 

SeCl44-3H2O=4HCl  +  H2SeO3. 

Corresponding  bromine  and  iodine  compounds  are  known,  SeBr4  and  SeI4. 


OXIDES  AND  OXYACIDS  OF  SELENIUM. 

Only  one  oxide  of  selenium  is  known,  namely,  selenium  dioxide,  SeO2, 
although  a  second  oxide  of  unknown  composition  is  believed  to  exist,  and  to 
constitute  the  peculiar  smelling  substance  whicli  is  always  formed  when 
selenium  is  burnt  in  the  air. 

Selenium  Dioxide  is  prepared  by  burning  selenium  in  a  stream  of  oxygen 
in  a  glass  tube  ;  the  element  burns  in  the  gas  with  a  blue  flame,  and  the  oxide 
condenses  upon  the  distant  portions  of  the  tube,  as  a  white  crystalline  deposit. 

Properties. — Selenium  dioxide  crystallises  in  long  white  prisms,  which  when 
heated  readily  sublime  without  passing  through  the  state  of  liquidity.  It  dis- 
..^Ives  in  water  and  gives  rise  to  selenious  acid. 

The  following  oxyacids  of  selenium  are  known — 

5      Selenious  acid,  H2SeO3,  corresponding  to  sulphurous  acid,  H2SO3. 
Selenic  acid,  H2SeO4,  corresponding  to  sulphuric  acid,  H2SO4. 
Selenosulphuric  HO   )  cr.       corresponding  (  thiosulphuric  HO  )  c~ 
acid  HSe  | b<Ja'  to  1         acid,          HS  /  b°2' 

Selenious  Acid,  H2SeO3,  is  obtained  as  a  white  crystalline  compound,  when 
the  dioxide  is  dissolved  in  hot  water,  and  the  solution  allowed  to  cool.  The 
acid  is  dibasic,  and  forms  both  acid  and  normal  selenites,  corresponding  to 
the  sulphites :  it  also  forms  a  series  of  so-called  superacid  salts,  containing  a 
molecule  of  the  acid  salt  combined  with  a  molecule  of  acid,  thus — 

HKSe03,H2SeO3. 

Selenic  Acid,  H2SeO4.—  This  acid  is  best  prepared  by  the  addition  of 
bromine  to  silver  selenite  suspended  in  water,  when  insoluble  silver  bromide  is 
formed  and  selenic  acid  is  It  ft  in  solution — 

Ag2SeO3  +  H2O  +  Br2=2AgBr  +  H2SeO4. 

The  solution  may  be  evaporated  by  heating  until  it  contains  94  per  cent,  of 
selenic  acid,  and  still  further  evaporated  in  vacuo  until  it  reaches  97.4  per  cent., 
when  its  specific  gravity  is  2.627.  When  heated  to  280°  it  decomposes  into 
selenium  dioxide,  water,  and  selenium. 

Properties. — Selenic  acid  in  its  most  concentrated  condition  is  a  colourless, 
strongly-acid  liquid,  which  mixes  with  water  with  the  development  of  con- 
siderable heat.  It  dissolves  iron  and  zinc  with  evolution  of  hydrogen  ;  and  when 
heated  dissolves  copper  with  formation  of  selenious  acid. 

The  selenates  closely  resemble  the  sulphates.      Barium  selenate,  like  the 


448  Inorganic  Chemistry 

sulphate,  is  quite  insoluble  in  water,  but  differs  from  that  compound  in  being 
converted  by  boiling  hydrochloric  acid  into  barium  selenite,  which  is  soluble. 

Selenium  also  forms  a  compound  with  oxygen  and  chlorine,  selenium  oxy- 
chloride,  or  selenyl  chloride,  SeOCl2,  corresponding  with  thionyl  chloride, 
SOC12. 

TELLURIUM. 

Symbol,  Te.     Atomic  weight*  =  127. 6. 

Occurrence. — In  the  free  state  small  quantities  of  this  element  have  been 
found  as  crystals,  consisting  of  almost  pure  tellurium.  In  combination  it  is 
met  with  in  a  few  rare  minerals,  such  as  tellurite  (TeO2),  and,  more  commonly, 
tetradymite  (Bi2Te3).  Some  specimens  of  pyrites  contain  small  quantities  of 
this  element,  hence  it  is  found  in  the  deposit  from  the  vitriol  chambers,  from 
which  selenium  is  obtained. 

Mode  of  Formation. —Tellurium  is  obtained  from  bismuth  telluride,  Bi2Te3, 
by  fusion  with  an  intimate  mixture  of  sodium  carbonate  and  carbon.  The 
mass  on  treatment  with  water  yields  a  solution  containing  a  mixture  of  sodium 
telluride  and  sodium  sulphide,  which  on  exposure  to  the  air  deposits  tellurium 
as  a  grey  powder.  The  element  is  purified  by  distillation  in  a  stream  of 
hydrogen. 

Properties. — Tellurium  is  a  bluish-white,  silver-like  solid,  possessing  metallic 
lustre.  It  conducts  heat  and  electricity,  although  badly,  and  is  very  brittle. 
Its  specific  gravity  is  6.26,  and  it  melts  at  452°.  When  melted  tellurium  is 
slowly  cooled,  it  forms  rhombohedral  crystals.  When  heated  in  the  air  it  burns 
with  a  blue  flame,  and  forms  tellurium  dioxide,  TeO2.  When  heated  in  a 
sealed  tube  with  hydrogen,  tellurium  sublimes  in  the  form  of  brilliant  prismatic 
crystals. 

Hydrogen  Telluride  (telluretted  hydrogen},  H2Te. — When  tellurium 
is  heated  in  hydrogen  the  elements  combine,  forming  hydrogen  telluride, 
which  exhibits  the  same  phenomenon  as  is  shown  by  hydrogen  selenide  of 
being  decomposed  by  heat,  and  depositing  the  element  as  a  crystalline 
sublimate. 

Hydrogen  telluride  is  obtained  by  the  action  of  hydrochloric  acid  upon  zinc 
telluride— 

ZnTe  +  2HC1  =  ZnCl2  +  K2Te. 

Properties. — Hydrogen  telluride  is  a  most  offensive  smelling  and  highly 
poisonous  gas.  It  behaves  like  sulphuretted  hydrogen  in  precipitating  metals 
from  solutions.  It  is  soluble  in  water,  and  the  solution  gradually  absorbs 
oxygen  and  deposits  tellurium. 

*  Various  numbers  have  been  obtained  by  different  observers  for  the 
atomic  weight  of  tellurium.  Some  of  these  numbers  are  higher  than  the  atomic 
weight  of  iodine,  which  would  make  it  impossible  to  give  to  tellurium  a  posi- 
tion between  antimony  (atomic  weight  =  120)  and  iodine  (atomic  weight  = 
126.97)  as  demanded  by  the  periodic  law.  Brauner,  who  has  spent  many 
years  investigating  this  point,  considers  that  hitherto  pure  tellurium  has  never 
been  obtained.  The  most  recent  determinations  give  the  number  127.6. 


Telluric  Acid  449 

COMPOUNDS  WITH  THE  HALOGENS. 

Two  chlorides  of  tellurium  are  known,  namely,  tellurium  dichloride,  TeCl2, 
and  tellurium  tetrachloride,  TeCl4.  It  will  be  noticed  that  the  composition  of 
the  dichloride  is  not  analogous  with  the  lower  chloride  of  either  selenium 
(Se2Cl2)  or  sulphur  (S2C12). 

Two  bromides,  TeBr2  and  TeBr4,  and  corresponding  iodides  are  known. 


OXIDES  AND  OXYACIDS  OF  TELLURIUM. 

Two  oxides  of  tellurium  are  known  with  certainty,  namely,  tellurium  dioxide, 
TeO2,  and  tellurium  trioxide,  TeO3,  which  give  rise  respectively  to  the  two 
acids,  tellurous  acid,  H2TeO3,  and  telluric  acid,  H2TeO4. 

Tellurous  Acid  is  obtained  by  pouring  a  solution  of  tellurium  in  nitric  acid 
into  an  excess  of  water.  The  acid  is  precipitated  as  a  white  amorphous 
powder.  When  strongly  heated  it  is  converted  into  the  dioxide  and  water. 

Tellurous  acid,  like  sulphurous  acid,  is  dibasic,  and  gives  rise  to  both  acid 
and  normal  salts  :  thus,  with  potassium  it  forms  hydrogen  potassium  tellurite, 
HKTeO3,  and  dipotassium  tellurite,  K2TeO3.  It  also  forms  superacid  salts 
such  as — 

Quadracid  potassium  tellurite      .         .         .     HKTeO3,H2TeO3. 
Potassium  tetratellurite        ....     K2TeO3,3TeO2. 

Telluric  Acid  is  prepared  by  fusing  either  tellurium  or  tellurium  dioxide 
with  a  mixture  of  potassium  nitrate  and  carbonate — 

Te:  +  K2C03  +  2KN03=2K2Te04  +  N2  +  CO. 

The  fused  mass,  after  solution  in  water,  is  mixed  with  a  solution  of  barium 
chloride,  which  precipitates  barium  tellurate  ;  this  is  then  decomposed  by  the 
addition  of  the  exact  amount  of  sulphuric  acid,  and  after  filtration  the  clear 
solution  deposits  crystals  of  telluric  acid,  H,TeO4,2H2O.  When  these  crystals 
are  heated  to  160°  the  water  is  expelled,  and  the  anhydrous  acid  in  the  form  of 
a  white  powder  is  left.  On  strongly  heating,  telluric  acid  decomposes  into 
water  and  tellurium  trioxide,  which  at  a  higher  temperature  splits  up  into  the 
dioxide  and  oxygen. 

Like  tellurous  acid,  telluric  acid  forms  not  only  normal  and  acid  salts,  but  a 
number  of  more  complex  superacid  salts — 

Normal  potassium  tellurate         .         .         .  K2TeO4,5H2O. 

Hydrogen  potassium  tellurate     .         .         .  HKTeO4. 

Quadracid  potassium  tellurate    .         .         .  K2TeO4,H,,TeO.1,3H2O. 

Potassium  tetratellurate      ....  K2TeO4,3H2TeO4,H2O. 


2F 


CHAPTER  III 
THE  ELEMENTS  OF  GROUP  V.  (FAMILY  B.) 


Nitrogen,  N      .        .     14.01 
Phosphorus,  P  .         .31.0 

Arsenic,  As       .         -75.0 


Antimony,  Sb     .         .     120.2 
Bismuth,  Bi        .        .     208 


IN  this  family  of  elements  we  have  a  gradual  transition  from  the 
non-metals  to  the  metals.  Nitrogen  and  phosphorus  may  be  con- 
sidered as  typical  non-metallic  elements,  both  as  regards  their 
physical  and  chemical  properties.  The  third  member,  arsenic, 
begins  to  exhibit  metalline  properties  ;  its  specific  gravity  is  more 
than  three  times  as  high  as  that  of  phosphorus,  and  it  possesses 
considerable  metallic  lustre ;  arsenic  is  called  a  metalloid  on  this 
account.  Antimony  is  still  more  metallic  in  its  character,  possess- 
ing most  of  the  physical  attributes  of  a  true  metal,  while  in  bismuth 
all  non-metallic  properties  cease  altogether  to  exist. 

All  these  elements  form  more  than  one  compound  with  oxygen, 
of  which  the  following  may  be  compared — 

N203  ;  (P203)2 ;  (As2O3)2  ;  Sb2O3  ;  Bi2O3. 
N204  j  P204  ;  Sb204  ;  Bi204. 

N205  ;   P205  ;       As205  ;        Sb2O5  ;   Bi2O5. 

The  oxides  (which  in  the  case  of  nitrogen  and  phosphorus  are 
strongly  acidic  in  their  nature,  combining  with  water  to  form  acids) 
gradually  become  less  and  less  acidic  and  more  basic  as  the  series 
is  traversed. 

Thus,  nitrogen  pentoxide,  N2O5,  unites  violently  with  water  to 
form  nitric  acid,  which  with  bases  yields  nitrates.  Antimony  pent- 
oxide  is  insoluble  in  water,  and  no  antimonic  acid  has  been  isolated, 
although  its  salts,  the  antimonates,  are  known.  The  oxides  of 
antimony,  on  the  other  hand,  begin  to  exhibit  basic  properties  and 
unite  with  acids,  forming  salts  in  which  the  antimony  functions  as 
the  base. 

450 


Phosphorus  451 

In  the  case  of  the  last  element  the  acidic  nature  of  the  oxides  is 
entirely  lost ;  no  bismuth  compounds  being  known  corresponding 
to  antimonates  or  arsenates,  while  these  oxides  unite  with  acids  in 
the  capacity  of  bases,  giving  rise  to  bismuth  salts. 

Four  of  the  elements  of  this  group  unite  with  hydrogen,  forming 
similarly  constituted  compounds,  NH3,  PH3,  AsH3,  SbH3. 

The  stability  of  these  compounds  gradually  decreases  as  we  pass 
from  nitrogen  to  antimony.  Antimony  hydride  has  never  been  ob- 
tained free  from  other  gases,  while  no  similar  bismuth  compound  is 
known.  Ammonia  is  alkaline  and  strongly  basic,  and  unites  readily 
with  acids  to  form  ammonium  salts.  Phosphorus  hydride  has  no 
alkaline  character,  and  is  only  feebly  basic.  It  combines,  however, 
with  the  halogen  acids  to  form  phosphonium  chloride,  bromide, 
and  iodide,  PH4C1,  PH4Br,  PH4I,  analogous  to  ammonium  chloride, 
bromide,  and  iodide.  The  hydrides  of  arsenic  and  antimony  ex- 
hibit no  basic  character.  All  the  elements  of  this  group  unite  with 
chlorine,  giving  rise  to  the  compounds — 

NC13,  PC13,  AsCl3,  SbCl3,  BiCl3, 

which  also  exhibit  a  gradation  in  their  properties ;  thus,  nitrogen 
trichloride  is  an  extremely  unstable  liquid,  exploding  with  extra- 
ordinary violence  upon  very  slight  causes,  while  the  analogous 
bismuth  compound  is  a  perfectly  stable  solid. 

The  boiling-points  of  these  compounds  show  a  gradual  increase 
with  the  increasing  atomic  weight  of  the  element  ;  thus,  nitrogen 
chloride  boils  at  71°,  phosphorus  trichloride  at  78°,  arsenic  tri- 
chloride at  130.2°,  and  antimony  trichloride  at  200°. 

The  elements  arsenic,  antimony,  and  bismuth  are  isomorphous, 
and  their  corresponding  compounds  are  also  isomorphous. 

The  first  member  of  this  family,  namely,  nitrogen,  has  been 
already  treated  in  Part  II.  as  one  of  the  four  typical  elements 
studied  in  that  section  of  the  book.  It  occupies  a  position  in 
relation  to  the  other  members  of  the  family  very  similar  to  that  of 
oxygen  towards  sulphur,  selenium,  and  tellurium. 


PHOSPHORUS. 

Symbol,  P.     Atomic  weight  =  31.0.     Vapour  density =62.0. 
Molecular  weight  =  124.0. 

History. — Phosphorus    was   first  discovered   by  the   alchemist: 
Brand  of  Hamburg  (1669),  who  obtained  it  by  distilling  a  mixture 


452  Inorganic  Chemistry 

of  sand  with  urine  which  had  been  evaporated  to  a  thick  syrup. 
The  process,  however,  was  kept  secret.  Robert  Boyle  (1680)  dis- 
covered the  process  of  obtaining  this  element,  but  the  method  was 
not  published  till  after  his  death.  Until  the  year  1771,  when 
Scheele  published  a  method  by  which  phosphorus  could  be  ob- 
tained from  bone  ash,  this  element  was  looked  upon  as  a  rare 
chemical  curiosity.  The  name  phosphorus  was  not  first  coined  for 
this  element :  it  had  been  in  previous  use  to  denote  various  sub- 
stances known  at  that  time,  which  had  the  property  of  glowing  in 
the  dark.  To  distinguish  the  element  it  was  called  Brand's  phos- 
phorus, or  English  phosphorus. 

Occurrence. — Phosphorus  is  not  found  in  nature  in  the  Tree 
state.*  In  combination  with  oxygen  and  metals,  as  phosphates, 
it  is  very  widely  distributed,  especially  as  calcium  phosphate. 
The  following  are  some  of  the  commonest  natural  phosphates — 

Sombrerite,  or  estramadurite       .      Ca3(PO4)2. 

Apatite 3Ca3(PO4)2,CaCl2. 

Wavellite 2A12(PO4),,A12(HO)6,9H2O. 

Calcium  phosphate  is  present  in  all  fertile  soils,  being  derived 
from  the  disintegration  of  rocks  :  the  presence  of  phosphates  in 
soil  has  been  shown  to  be  essential  to  the  growth  of  plants.  From 
the  vegetable  it  passes  into  the  animal  kingdom,  where  it  is  chiefly 
present  in  the  urine,  brain,  and  bones.  Bones  contain  about  60 
per  cent,  of  calcium  phosphate,  to  which  they  entirely  owe  their 
rigidity. 

Mode  Of  Formation.—  Manufacture.  The  chief  source  of 
phosphorus  is  bone  ash,  a  material  obtained  by  burning  bones, 
and  which  consists  of  nearly  pure  calcium  phosphate,  Ca3(PO4)2. 
Other  varieties  of  calcium  phosphate,  such  as  sombrerite  and 
apatite,  are  also  employed,  as  well  as  phosphates  of  other  metals, 
such  as  the  Redonda  phosphates,  which  consist  of  phosphates  of 
iron  and  alumina.  The  bone  ash,  in  fine  powder,  is  first  decom- 
posed by  means  of  sulphuric  acid,  specific  gravity  1.5  to  1.6.  This 
operation  is  performed  in  large  circular  wooden  vessels,  resem- 
bling a  brewer's  "mash  tun,"  provided  with  an  agitator,  and  into 
which  high  pressure  steam  can  be  driven.  Finely-ground  bone  ash 
and  sulphuric  acid,  in  charges  of  a  few  cwts.  at  a  time,  are  alter- 
nately stirred  into  the  decomposer,  until  from  four  to  five  tons  of 

*  Farrington  (Am.  Jour.  Scie?ice,  vol.  xv. ,  1903)  records  having  discovered 
small  quantities  cf  free  phosohcrus  enclosed,  or  occluded  in  a  meteorite. 


Phosphorus 


453 


phosphate  have  been  introduced,  with  sufficient  acid  to  convert  the 
whole  of  the  lime  into  calcium  sulphate,  according  to  the  equa- 
tion — 


The  contents  of  the  decomposer  are  next  run  out  into  filtering- 
tanks,  and  the  phosphoric  acid  is  then  concentrated  to  a  syrup,  in 
large  lead-lined  pans  through  which  steam-coils  pass,  the  liquor 
being  constantly  agitated  by  a  mechanical  stirrer. 

The  concentrated  liquor  is  next  mixed,  either  with  sawdust,  or 
with  coarsely-ground  charcoal,  or  coke,  and  the  mixture  com- 
pletely dried  by  being  heated  in  a  cast-iron  pot,  or  in  a  muffle,  to  a 


dull  red  heat.  During  this  process  the  tribasic  phosphoric  acid 
(or  orthophosphoric  acid),  H3PO4,  is  converted  by  loss  of  water 
into  metaphosphoric  acid,  HPO3  — 

H3P04=H2O  +  HPO3. 

The  charred  mixture  is  then  distilled  in  bottle-shaped  retorts  of 
Stourbridge  clay,  about  3  feet  long,  and  having  an  internal  diameter 
of  8  inches.  A  number  of  these  retorts,  usually  twenty-four,  are 
arranged  in  two  tiers,  in  a  galley  furnace,  as  seen  in  section  in  Fig. 
1 1 8.  The  empty  retorts  are  first  gradually  raised  to  a  bright  red 
heat,  and  a  charge  of  the  mixture  is  then  quickly  introduced.  Bent 


454 


Inorganic  Chemistry 


pieces  of  2-inch  malleable  iron  pipe  are  then  luted  into  the 
mouths  of  the  retorts  connecting  them  with  the  pipes,  D  D'. 
These  pipes  dip  into  troughs  of  water,  E  E',  which  run  along  the 
entire  length  of  the  furnace,  and  in  which  the  phosphorus  con- 
denses. The  temperature  of  the  furnace  is  then  raised  to  a  white 
heat,  when  decomposition  of  the  metaphosphoric  acid  commences, 
and  phosphorus  begins  to  distil  over.  The  process  is  continued 
for  about  sixteen  hours.  The  change  that  goes  on  is  mainly 
represented  by  the  following  equation  — 


4HP03+.12C  = 
The  crude  product,  which  is  usually  dark  red  or  black  in  appear- 


FIG.  119. 

ance,  is  first  melted  under  hot  water  and  thoroughly  stirred,  in 
order  to  allow  the  greater  part  of  the  rougher  suspended  matters  to 
rise  to  the  surface.  The  mass  is  then  allowed  to  resolidify.  The 
exact  processes  by  which  phosphorus  is  further  purified  on  a  manu- 
facturing scale  are  guarded  as  trade  secrets  ;  one  method  that  has 
been  in  use  consists  in  treating  the  phosphorus  while  melted  under 
water  with  a  mixture  of  potassium  dichromate  and  sulphuric  acid, 
whereby  some  of  the  impurities  are  oxidised  and  others  are  caused 


Phosphorus  45  5 

to  rise  to  the  surface  as  a  scum,  leaving  the  phosphorus  as  a  clear 
liquid  beneath. 

Since  the  advent  of  the  electric  furnace,  phosphorus  is  now  being 
manufactured  direct  from  calcium  phosphate  by  a  process  which 
threatens  to  entirely  supersede  the  method  of  distillation  already 
described.  The  calcium  phosphate  is  mixed  with  carbon  and 
simply  heated  in  the  electric  furnace.  At  the  high  temperature  of 
the  electric  arc  the  calcium  phosphate  is  decomposed,  the  calcium 
uniting  with  the  carbon  to  form  calcium  carbide,  while  the  phos- 
phorus in  the  state  of  vapour  escapes  along  with  carbon  monoxide 
by  the  pipe  P  (Fig.  119),  and  is  condensed  in  suitable  condensers — 

Ca3(PO4)2  +  14C  =  3CaC2+ 2P  +  SCO. 
The  molten  calcium  carbide  is  tapped  off  from  the  furnace  from 


FIG.  120. 

time  to  time  as  fresh  charges  of  phosphate  and  carbon  are 
introduced. 

Phosphorus  usually  comes  into  commerce  either  in  the  form  of 
wedges  or  as  sticks.  The  operation  of  casting  the  phosphorus  into 
sticks  is  performed  beneath  water.  A  quantity  of  phosphorus 
beneath  a  shallow  layer  of  water  is  placed  in  the  vessel  C  (Fig.  120), 
which  is  contained  in  a  tank  of  water  through  which  a  steam-coil 
passes.  Connected  to  the  phosphorus  reservoir  is  a  glass  tube,  G, 
which  passes  into  a  second  shallow  tank  of  cold  water.  On  open- 
ing the  cock  D,  the  liquid  phosphorus  flows  into  the  cold  glass  tube 
where  it  congeals,  and  it  may  then  be  drawn  through  as  a  continuous 
rod  of  phosphorus  if  care  be  taken  not  to  draw  it  out  faster  than  it 
solidifies.  It  is  the  custom  to  adopt  a  uniform  length  and  thick- 
ness of  stick,  namely,  7^  inches  long  and  \  inch  diameter.  Nine 
such  sticks  weigh  I  Ib. 

Properties. — When  freshly  prepared  and  kept  in  the  dark, 
phosphorus  is  a  translucent,  almost  colourless,  wax-like  solid. 
Even  in  the  dark  it  soon  loses  its  transparency  and  becomes 


456  Inorganic  Chemistry 

coated  with  an  opaque  white  film  ;  while  if  exposed  to  the  light 
the  film  that  forms  becomes  first  yellow,  then  brown,  and  in  time 
the  phosphorus  assumes  a  red  and  even  a  black  colour  throughout 
its  entire  mass.  Its  specific  gravity  at  16°  is  1.82.  At  o°  phos- 
phorus becomes  moderately  brittle,  and  a  stick  of  it  may  be  readily 
snapped,  when  its  crystalline  character  will  be  seen.  At  1 5°  it 
becomes  soft,  and  may  be  cut  with  a  knife  like  wax.  Phosphorus 
melts  under  water  at  43.3°,  and  the  liquid  exhibits  the  property  of 
suspended  solidification.  If  the  melted  material,  which  has  been 
cooled  below  its  solidifying  point,  be  touched  with  a  fragment  of 
phosphorus  upon  the  end  of  a  capillary  glass  tube,  the  mass 
instantly  congeals  with  rise  of  temperature.* 

Phosphorus  contained  in  a  closed  vessel  without  water  melts 
at  as  low  a  temperature  as  30°,!  and  when  heated  in  air  to  34°  it 
takes  fire.  At  a  temperature  of  269°  phosphorus  boils  and  forms 
a  colourless  vapour. 

Phosphorus  is  volatile  at  ordinary  temperatures  :  if  a  small 
quantity  of  phosphorus  be  sealed  in  a  vacuous  tube,  and  the  tube 
be  placed  in  the  dark,  the  phosphorus  will  slowly  vaporise  ;  and  if 
one  end  of  the  tube  be  kept  slightly  cooler  than  the  rest,  the  phos- 
phorus will  sublime  upon  that  part  in  the  form  of  brilliant,  colour- 
less^ and  highly  refracting  rhombic  crystals,  which  retain  their 
beauty  as  long  as  they  are  kept  in  the  dark.  The  density  of 
the  vapour  of  phosphorus  is  62.0,  giving  a  molecular  weight  of 
124.0,  which  is  four  times  the  atomic  weight,  showing  that  the 
molecule  of  phosphorus  contains  four  atoms.  Even  at  temperatures 
as  high  as  1040°  these  tetratomic  molecules  are  stable,  but  it  has 
been  shown  that  at  high  temperatures  dissociation  begins  to  take 
place. 

On  account  of  its  ready  inflammability,  phosphorus  is  always 
preserved  under  water,  which  exerts  practically  no  solvent  action 
upon  it.  It  is  extremely  soluble  in  carbon  disulphide,  i  part  of  this 
liquid  dissolving  9.26  parts  of  phosphorus.  On  evaporation,  the 
element  is  deposited  in  the  form  of  colourless  crystals.  Phosphorus 
is  also  soluble,  but  to  a  less  extent,  in  chloroform,  benzene,  turpen- 
tine, alcohol,  olive  oil,  and  many  other  solvents.  A  solution  of 
phosphorus  in  carbon  disulphide,  when  allowed  to  evaporate  upon 
a  piece  of  blotting-paper,  leaves  the  element  in  so  finely  divided  a 
condition,  that  its  rapid  oxidation  almost  immediately  raises  the 

*  See  "Chemical  Lecture  Experiments,"  new  ed.f  Nos.  528,  529. 
f  Readman. 


Phosphorus  457 

temperature  to  the  ignition  point  of  the  phosphorus,  when  it 
takes  fire. 

On  exposure  to  moist  air  in  the  dark,  phosphorus  appears  faintly 
luminous,  emitting  a  pale  greenish-white  light,  and  at  the  same 
time  evolving  white  fumes  which  possess  an  unpleasant,  garlic-like 
smell,  and  are  poisonous.  These  fumes  consist  mainly  of  phos- 
phorus oxide,  P4O6,  and  the  glowing  of  the  phosphorus  is  the 
result  of  its  oxidation  ;  phosphorus  does  not  glow  when  placed 
in  an  inert  gas  which  is  perfectly  free  from  admixed  oxygen, 
although  the  presence  of  very  small  traces  of  free  oxygen  in  such  a 
gas  is  sufficient  to  cause  the  phosphorescence.  At  a  few  de- 
grees below  o°,  phosphorus  ceases  to  glow  in  the  air.  Although 
the  glowing  is  due  to  oxidation,  phosphorus  does  not  appear 
luminous  in  pure  oxygen  at  temperatures  below  about  15°.  If, 
therefore,  a  stick  of  phosphorus  which  is  glowing  in  the  air, 
be  immersed  in  a  jar  of  oxygen,  its  phosphorescence  is  at  once 
stopped.  If,  however,  the  oxygen  be  slightly  rarefied,  the  phos- 
phorus again  becomes  luminous.  Similarly,  the  phosphorescence 
that  is  exhibited  in  air  is  stopped  if  the  air  be  compressed.*  The 
glow  of  phosphorus  is  believed  to  be  associated  with  the  formation 
of  ozone,  for  the  presence  in  the  air  of  traces  of  such  gases  and 
vapours  as  ethylene,  turpentine,  or  ether,  which  are  known  to 
possess  the  power  of  destroying  ozone,  at  once  stops  the  glowing 
of  a  stick  of  phosphorus. 

Phosphorus  is  incapable  of  uniting  with  oxygen  if  the  gas  be 
perfectly  pure  and  free  from  aqueous  vapour.  It  has  been  shown 
that  in  oxygen  which  has  been  dried  by  prolonged  exposure  to  the 
desiccating  action  of  phosphorus  pentoxide,  phosphorus  may  not 
only  be  melted,  but  even  distilled,  without  any  combination  with 
the  oxygen  taking  place. 

If  water,  beneath  which  is  a  small  quantity  of  melted  phosphorus, 
be  boiled,  the  phosphorus  vaporises  with  the  steam,  and  renders 
the  steam  luminous  :  use  is  made  of  this  property,  as  a  means 
of  detecting  free  phosphorus,  in  toxicological  analysis. 

Phosphorus  is  a  powerfully  poisonous  substance  ;  in  large  doses 
it  causes  death  in  a  few  hours,  in  smaller  quantities  it  produces 
stomachic  pains  and  sickness,  usually  ending  in  convulsion, 
Persons  constantly  exposed  to  the  vapours  arising  from  the  hand- 
ling of  phosphorus,  either  in  its  manufacture  or  in  the  manufacture 

*  "  Chemical  Lecture  Experiments,"  new  ed.,  Nos.  530  to  534. 


458 


Inorganic  Chemistry 


of  matches,  are  very  liable  to  suffer  from  caries  of  the  bones  of  the 
jaw  and  nose  ;  it  is  believed  that  this  injurious  effect  is  caused  by 
the  white  fumes  which  are  the  product  of  oxidation,  and  not  by 
the  actual  vapour  of  phosphorus. 

Red  Phosphorus.— -When  phosphorus  is  heated  to  a  tempera- 
ture between  240°  and  250°,  out  of  contact  with  air,  it  passes  into 
an  allotropic  modification.  The  same  transformation  takes  place 
when  phosphorus  is  heated  to  200°  with  an  extremely  small  pro- 
portion of  iodine. 

Red  phosphorus  is  manufactured  by  heating  ordinary  phosphorus 


FIG.  121. 

in  a  cast-iron  pot,  provided  with  a  cover,  through  which  passes  a 
short  open  pipe,  D  (Fig.  121).  The  pot  is  carefully  and  uniformly 
heated  to  between  240°  and  250°,  as  indicated  by  the  thermometers 
C  C',  which  are  encased  in  metal  tubes,  to  prevent  the  phosphorus 
from  attacking  the  glass.  A  small  quantity  of  the  phosphorus 
becomes  oxidised  by  the  air  within  the  vessel,  but  after  this  atmos- 
pheric oxygen  is  used  up,  no  further  oxidation  takes  place.  If  the 
temperature  be  allowed  to  rise  above  260°,  the  red  phosphorus  is 


Red  Phosphorus  459 

reconverted  into  the  ordinary  modification,  and  with  the  evolution 
of  so  much  heat,  that  unless  the  open  tube  be  provided,  as  a  safety- 
valve,  the  iron  vessel  is  liable  to  burst.  The  material  that  is 
obtained  at  the  end  of  the  operation  is  in  the  form  of  hard,  solid 
lumps,  which  still  contain  a  certain  amount  of  the  unchanged 
phosphorus  mixed  with  them.  It  is  first  ground  to  powder  beneath 
water,  and  then  boiled  with  a  solution  of  sodium  hydroxide  (caustic 
soda),  to  remove  the  ordinary  phosphorus,  and  finally  washed  and 
dried. 

Properties. — Red  phosphorus,  as  usually  sent  into  commerce, 
is  a  chocolate-red  powder,  having  a  specific  gravity  of  2.25.  It 
is  not  luminous  in  the  dark,  and  has  no  taste  or  smell.  It  is 
not  poisonous,  and  when  taken  into  the  system  is  excreted  un- 
changed. It  is  not  soluble  in  carbon  disulphide,  or  in  any  of  the 
solvents  which  dissolve  ordinary  phosphorus.  Red  phosphorus  is 
unaffected  by  exposure  to  dry  air  or  oxygen,  but  in  the  presence 
of  moisture  it  is  very  slowly  oxidised.  If  red  phosphorus  which 
has  been  perfectly  freed  from  ordinary  phosphorus,  and  carefully 
washed  and  dried,  be  exposed  to  air  and  moisture,  it  is  found  after 
the  lapse  of  some  time  to  have  become  acid,  owing  to  slight  oxida- 
tion into  phosphoric  acid.  When  heated  in  contact  with  air,  red 
phosphorus  does  not  ignite  below  a  temperature  of  240°.  Red 
phosphorus  may  be  obtained  in  the  form  of  rhombohedral  crystals 
by  heating  the  substance  under  pressure  to  a  temperature  of  580°. 

The  chief  use  of  phosphorus  is  in  the  manufacture  of  matches. 
When  ordinary  phosphorus  is  employed,  the  bundles  of  wooden 
splints  are  first  tipped  with  melted  paraffin  wax,  and  afterwards 
dipped  into  a  paste,  made  of  an  emulsion  of  phosphorus,  chlorate 
of  potash,  and  glue.  Matches  so  made  ignite  when  rubbed  upon 
any  rough  surface  ;  the.  paraffin  (which  is  sometimes  replaced  by 
sulphur)  serving  to  transmit  the  combustion  from  the  phosphorus 
to  the  wood.  Since  the  discovery  of  red  phosphorus,  and  its  non- 
injurious  properties,  the  old  phosphorus  match  has  been  largely 
superseded  by  the  so-called  safety  matches.  In  these  matches  the 
splints  are  tipped  with  a  mixture  of  potassium  chlorate,  potassium 
dichromate,  red  lead,  and  antimony  sulphide,  and  they  are  ignited 
by  being  rubbed  upon  a  prepared  surface  consisting  of  a  mixture 
of  antimony  sulphide  and  red  phosphorus.  Although  these  matches 
will  not  ignite  by  ordinary  friction  upon  any  but  the  specially 
prepared  surface,  they  may  be  inflamed  by  being  swiftly  drawn 
along  a  sheet  of  ground  glass  or  strip  of  linoleum. 


460  Inorganic  Chemistry 

COMPOUNDS    OF  PHOSPHORUS    WITH  HYDROGEN. 

Three   compounds    of  phosphorus   and   hydrogen   are   known, 
namely — 

PH3  (gaseous)  ;  P2H4  (liquid)  ;.  and  (P4H2)3  (solid). 


GASEOUS  HYDROGEN  PHOSPHIDE  (Phosphorated  Hydrogen  : 
Phosphine). 

Formula,  PH3.     Molecular  weight  =  34.  03.     Density=  17.015. 

Modes  Of  Formation.—  (i.)  This  substance  is  formed  when  red 
phosphorus  is  gently  heated  in  a  stream  of  hydrogen. 

(2.)  It  may  be  prepared  by  boiling  phosphorus  with  a  solution  of 
potassium  or  sodium  hydroxide  — 

3NaHO  +  4P  +  3H2O  =  3NaH2PO2  +  PH3. 

In  this  reaction  a  small  quantity  of  the  liquid  hydride  (P2H4)  is 
simultaneously  formed,  which  imparts  to  the  gas  the  property  of 
spontaneous  inflammability.  It  also  contains  a  certain  quantity 
of  free  hydrogen,  produced  by  the  action  of  the  caustic  alkali  upon 
the  sodium  hypophosphite,  thus  — 


To  obtain  the  gas  by  this  method,  a  quantity  of  a  strong  solution 
of  caustic  soda,  and  a  few  fragments  of  phosphorus,  are  placed  in 
a  flask,  fitted  as  shown  in  Fig.  122.  A  stream  of  coal  gas  is  passed 
through  the  apparatus,  in  order  to  displace  the  air,  and  the  solution 
is  gently  heated.  Hydrogen  phosphide  is  readily  disengaged,  and 
as  each  bubble  escapes  into  the  air,  it  bursts  into  flame,  and  forms 
a  vortex  ring  of  white  smoke  of  phosphoric  acid. 

If  alcoholic  potash  be  substituted  for  the  aqueous  solution  in 
this  reaction,  the  liquid  hydrogen  phosphide  is  dissolved  in  the 
alcohol,  and  the  gas  which  is  evolved  IB  therefore  not  spontaneously 
inflammable.* 

(3.)  Hydrogen  phosphide  is  also  produced  by  the  action  of 
water  upon  calcium  phosphide  — 


*  See  "  Chemical  Lecture  Experiments,"  new  ed.  ,  No.  545. 


Hydrogen  Phosphide 


461 


A  secondary  reaction,  by  which  liquid  hydrogen  phosphide  is 
formed,  goes  on  simultaneously  — 


The  gas,  therefore,  that  is  evolved  is  spontaneously  inflammable. 

(4.)  Pure  gaseous  hydrogen  phosphide  may  be  prepared  by  the 
action  of  potassium  hydroxide  upon  phosphonium  iodide  — 


Properties.  —  Gaseous  hydrogen  phosphide,  or  phosphine,  is  a 
colourless  gas,  having  an  offensive  smell  resembling  rotten  fish. 
It  is  not  spontaneously  inflammable,  but  its  ignition  temperature  is 


FIG.  122. 

below  100°  C.  (see  p.  329).    The  gas  burns  with  a  brightly  luminous 
flame,  producing  water  and  metaphosphoric  acid  — 

PH3  +  2O2=HPO3+H2O. 

When  burnt  in  oxygen  the"  flame  is  extremely  dazzling. 

The  gas  is  not  acted  upon  by  oxygen  at  ordinary  temperatures 
and  pressures,  but  if  a  mixture  of  these  gases  be  suddenly  rarefied, 
combination  at  once  takes  place  with  explosion.  Hydrogen  phos- 
phide is  decomposed  by  chlorine  or  bromine,  a  jet  of  the  gas 
spontaneously  igniting  when  introduced  into  chlorine  or  the  vapour 
of  bromine  — 


The  gas  is  also  decomposed  by  iodine,   but  in    this  case    the 
action  is  less  energetic,  and  a  portion  of  the  hydriodic  acid  which 


462  Inorganic  Chemistry 

is  produced  unites  with  the  phosphine  and  forms  phosphonium 
iodide,  thus — 

(i.)  PH3  +  3I,  =  PI3  +  3HI. 

(2.)  PH3+HI  =  PH4I. 

Phosphine  is  a  highly  poisonous  gas,  and  the  inhalation  of  even 
small  quantities  of  it  is  attended  with  injurious  effects.  It  is 
slightly  soluble  in  water,  and  imparts  its  own  smell  and  an  un- 
pleasant taste  to  the  liquid.  The  solution  decomposes  after  a 
short  time,  especially  in  the  light,  and  deposits  red  phosphorus. 

Hydrogen  phosphide  has  no  action  upon  either  litmus  or  tur- 
meric paper,  but  it  resembles  its  nitrogen  analogue,  ammonia, 
in  combining  with  hydrochloric,  hydrobromic,  andhydriodic  acids, 
forming  respectively  phosphonium  chloride,  bromide,  and  iodide. 

Phosphonium  Chloride,  PH4C1.— When  a  mixture  of  phos- 
phine and  gaseous  hydrochloric  acid  is  passed  through  a  tube 
immersed  in  a  freezing-mixture,  the  gases  unite  and  form  a  white 
crystalline  incrustation  upon  the  tube.  If  the  tube  be  afterwards 
sealed  up,  the  compound  may  be  sublimed  from  one  part  of  the 
tube  to  another,  when  it  crystallises  in  large,  brilliant,  transparent 
cubes.  If  the  tube  be  opened,  the  compound  rapidly  dissociates 
into  its  two  generators.  This  compound  may  also  be  obtained  by 
subjecting  a  mixture  of  the  two  gases  to  pressure.  Under  a  pres- 
sure of  about  eighteen  atmospheres  at  the  ordinary  temperature, 
crystals  of  phosphonium  chloride  are  deposited  ;  and  as  the  pres- 
sure is  released  the  crystals  gradually  dissociate  again. 

Phosphonium  Bromide,  PH4Br.— Hydrogen  phosphide  com- 
bines with  hydrobromic  acid  at  ordinary  temperatures  and  pres- 
sures, but  as  the  compound  begins  to  dissociate  at  the  ordinary 
temperature,  the  combination  is  only  completely  brought  about  by 
cooling  the  gases.  Phosphonium  bromide  may  be  readily  pre- 
pared by  passing  the  two  gases  into  a  flask  immersed  in  a  mode- 
rate freezing-mixture.  The  salt  may  be  obtained  in  the  form  of 
large  transparent  cubical  crystals  by  sublimation  in  a  sealed 
vessel. 

Phosphonium  Iodide,  PH4L— This  compound  may  be  obtained 
by  a  method  similar  to  that  given  for  the  preparation  of  the  bro- 
mide. It  is  also  produced  when  hydrogen  phosphide  is  passed 
over  iodine,  as  already  mentioned.  It  is  most  readily  prepared  by 
the  action  of  water  upon  a  mixture  of  phosphorus  and  iodine.  For 
this  purpose  ten  parts  of  phosphorus  are  dissolved  in  carbon  disul- 


Liquid  Hydrogen  Phosphide  463 

phide  in  a  tubulated  retort,  to  which  seventeen  parts  of  iodine  are 
gradually  added,  the  retort  being  kept  cold.  The  carbon  disul- 
phide  is  then  distilled  off  from  a  water-bath,  a  stream  of  carbon 
dioxide  being  passed  through  the  apparatus  towards  the  end  of 
the  distillation  to  assist  in  expelling  the  last  traces  of  the  disul- 
phide. 

Six  parts  of  water  are  then  gradually  introduced  from  a  dropping 
funnel,  when  a  brisk  action  takes  place,  and  the  phosphonium 
iodide  produced  is  volatilised,  and  may  be  condensed  in  a  long 
wide  glass  tube  connected  to  the  retort.  Hydriodic  acid  is  at  the 
same  time  formed  — 


The  phosphonium  iodide  condenses  in  the  form  of  brilliant 
quadratic  prisms. 

Liquid  Hydrogen  Phosphide,  I\H4.  —  This  compound  is 
obtained  in  small  quantities  when  phosphorus  is  boiled  with  a 


FIG.  123. 

solution  of  caustic  soda.  It  is  obtained  in  large  quantities  by  the 
decomposition  of  calcium  phosphide  with  water  by  the  reaction 
already  mentioned.  In  order  to  collect  the  compound  a  quantity 
of  calcium  phosphide  is  introduced  into  a  flask  provided  with  a 
dropping  funnel  and  exit  tube.  After  displacing  the  air  from  the 
apparatus  by  an  inert  gas  water  is  gradually  introduced  from  the 
funnel,  and  the  products  of  the  reaction,  after  passing  through  a 
small  empty  tube  where  water  is  arrested,  are  passed  through  a 
U-tube  immersed  in  a  freezing-mixture,  where  the  liquid  hydrogen 
phosphide  condenses. 

Properties. — Liquid  hydrogen  phosphide  is  a  colourless,  highly 
refracting,  and  spontaneously  inflammable  liquid.  On  exposure 
to  light  it  is  quickly  decomposed  into  the  gaseous  and  the  solid 
hydrides  of  phosphorus — 


464  Inorganic  Chemistry 

The  formation  of  a  spontaneously  inflammable  gas  by  the  action 
of  water  upon  calcium  phosphide  has  found  a  practical  application 
in  the  marine  appliance  known  as  "  Holmes'  signal."  This  con- 
sists of  a  tin  canister  filled  with  lumps  of  calcium  phosphide.  A 
metal  tube,  closed  at  the  bottom  with  a  piece  of  block  tin,  enters  the 
canister  from  below,  and  a  short  cone  of  the  same  soft  metal  is 
soldered  upon  the  top.  When  the  signal  is  to  be  used  it  is  securely 
fixed  into  a  wooden  float  (Fig.  123).  The  cone  is  cut  off  and  a 
hole  punctured  through  the  bottom  of  the  tube  below,  and  the 
apparatus  thrown  into  the  sea.  The  hydrogen  phosphide  spon- 
taneously ignites  and  burns  with  a  large  brilliant  flame  from  the 
top  of  the  tin,  illuminating  a  considerable  area. 

Solid  Hydrogen  Phosphide,  (P4H2)3  or  P12HG,  is  a  yellow 
powder,  obtained,  as  already  mentioned,  by  the  spontaneous 
decomposition  of  the  liquid  compound.  Recent  determinations* 
of  its  molecular  weight  prove  that  its  molecular  composition  is 
expressed  by  the  formula  P12H6. 

COMPOUNDS  OF  PHOSPHORUS  WITH  THE  HALOGENS. 
Phosphorus  combines  with  all  the  halogen  elements,  forming  the 
following  compounds  — 

PF3  PC13  PBr3  PI3. 

PF5  PC15  PBr5  P2I4. 

Phosphorus  Trifluoride,  PF3,  is  obtained  by  the  action  of 
arsenic  trifluoride  upon  phosphorus  trichloride  — 


It  is  more  conveniently  prepared  by  gently  heating  a  mixture  of 
zinc  fluoride  and  phosphorus  tribromide  — 


Properties.  —  Phosphorus  trifluoride  is  a  colourless,  pungent- 
smelling  gas.  It  has  no  action  upon  glass  in  the  cold,  but  when 
heated  it  forms  silicon  fluoride  and  phosphorus.  It  is  moderately 
soluble  in  water.  Phosphorus  trifluoride  unites  directly  with 
bromine,  forming  the  compound  PF3Br2. 

Phosphorus  Pentafluoride,  PF5.—  This  compound  is  formed 
when  phosphorus  burns  in  fluorine.  It  is  best  prepared  by  the 
action  of  arsenic  trifluoride  upon  phosphorus  pentachloride  — 


*  Schenck  and  Buck,  Berichte,  1904. 


Phosphorus  Pentachloride  465 

Properties.  —  Phosphorus  pentafluoride  is  a  heavy,  colourless 
gas,  which  fumes  strongly  in  moist  air,  being  decomposed  by  water 
into  hydrofluoric  and  phosphoric  acids— 


Owing  to  this  decomposition  it  has  a  pungent  and  irritating 
effect  upon  the  mucous  membrane. 

It  is  not  acted  upon  by  oxygen,  but  unites  directly  with  dry 
gaseous  ammonia,  forming  a  white  solid  compound  having  the 
composition  2PF5,5NH3. 

Phosphorus  pentafluoride  is  an  extremely  stable  compound,  being 
capable  of  withstanding  a  very  high  temperature  without  dissocia- 
tion. On  this  account  it  is  of  special  interest,  as  affording  an 
example  of  a  compound  in  which  phosphorus  is  united  to  five 
monovalent  atoms  to  form  a  stable  substance.  The  corresponding 
chlorine  and  bromine  compounds  readily  dissociate,  when  heated, 
into  compounds  containing  trivalent  phosphorus  and  the  free 
halogen. 

Phosphorus  Trichloride,  PC13.—  This  compound  is  prepared 
by  passing  dry  chlorine  over  red  phosphorus,  gently  heated  in  a 
tubulated  retort.  The  two  elements  readily  combine,  and  the 
volatile  trichloride,  mixed  with  more  or  less  of  the  pentachloride, 
distils  off,  and  is  collected  in  a  well-cooled  receiver.  The  product 
is  freed  from  the  higher  chloride  by  redistillation  over  ordinary 
phosphorus. 

Properties.  —  Phosphorus  trichloride  is  a  colourless,  mobile 
liquid,  which  boils  at  75.95°.  It  has  a  pungent  smell,  and  fumes 
strongly  in  moist  air.  Water  at  once  decomposes  it  into  hydro- 
chloric and  phosphorous  acids  — 


Phosphorus  Pentaehloride,  PC15.—  This  compound  is  formed 
when  phosphorus  burns  in  excess  of  chlorine.  It  is  prepared  by 
the  action  of  chlorine  upon  the  trichloride.  Dry  chlorine  is  passed 
on  to  the  surface  of  a  quantity  of  the  trichloride,  contained  in  a 
flask  which  is  kept  cool.  The  absorption  of  the  chlorine  is  attended 
with  considerable  rise  of  temperature,  and  the  contents  of  the  flask 
rapidly  become  converted  into  a  dry,  pale-yellow  solid. 

Phosphorus  pentachloride  is  conveniently  obtained  by  passing 

2  G 


466  Inorganic  Chemistry 

chlorine  through  a  solution  of  phosphorus  in  carbon  di  sulphide, 
the  solution  being  kept  cold. 

Properties.  —  Phosphorus  pentachloride  is  a  yellowish-white, 
crystalline  solid,  having  a  pungent  and  irritating  odour.  It  fumes 
strongly  in  contact  with  moist  air,  being  decomposed  by  moisture 
into  hydrochloric  acid  and  phosphorus  oxychloride— 

PC15  +  H,0  =  2HC1  +  POC13. 

With  excess  of  water,  both  phosphorus  oxychloride  and  phos- 
phorus pentachloride  dissolve  with  evolution  of  heat,  forming 
hydrochloric  and  phosphoric  acids— 


POC13  +  3H2O  =  H3PO4  +  3HC1. 
PC15     +4H2O  =  H3PO4  +  5HC1. 

Phosphorus  pentachloride  readily  sublimes,  without  melting,  at 
a  temperature  below  that  of  boiling  water.  It  can  only  be  melted 
by  being  heated  under  pressure  to  a  temperature  of  148°. 

As  the  vapour  of  phosphorus  pentachloride  is  heated,  the  com- 
pound dissociates  into  phosphorus  trichloride  and  free  chlorine. 
At  300°  this  dissociation  is  complete,  and  the  vapour  consists  of 
equal  molecules  of  the  trichloride  and  chlorine.  The  gradual 
breaking  down  of  the  molecules  of  pentachloride  is  seen  from  the 
following  table,  which  gives  the  densities  of  the  gas  at  different 
temperatures  — 

Temperatures      182°  200°  250°  300° 

Density      .     .      72.5  69.2  57.0  52.06 

At  300°  it  consists  of  molecules  of  PC13  (molecular  weight  = 
1  37-35)  and  molecules  of  chlorine  (molecular  weight  =  70.90)  in 
equal  numbers,  which  theoretically  gives  the  molecular  weight  — 


Phosphorus  pentachloride  is  an  important  chemical  reagent,  in- 
asmuch as  by  its  action  upon  oxyacids,  both  inorganic  and  organic. 
the  (HO)  group  in  the  acid  can  be  replaced  by  chlorine.  Thus 
with  sulphuric  acid,  chlorosulphuric  acid  is  formed  — 


Phosphorus  Pentachloride  467 

With  acetic  acid  it  yields  acetyl  chloride  — 


} 


CH 


It  also  effects  the  replacement  of  (HO)  by  chlorine,  in  alcohols. 
Thus,  with  ethyl  alcohol  (spirits  of  wine)  it  forms  ethyl  chloride  — 

C  Cl  |  CH2+POC13+HC1. 

Phosphorus  Triforomide,  PBr3,  is  best  prepared  by  dropping  bromine  upon 
an  excess  of  red  phosphorus.  It  forms  a  colourless  pungent-smelling  liquid, 
which  boils  at  172.9°. 

Phosphorus  Pentabromide,  PBr5,  is  prepared  by  adding  bromine  to  the 
tribromide.  It  is  a  yellow  solid,  which  melts  to  a  reddish  liquid.  It  is  very 
unstable,  being  dissociated  below  100°  into  its  generators,  the  tribromide  and 
bromine. 

Di  phosphorus  Tetriodide  (phosphorus  di-iodide],  P2I4.—  This  substance  is 
prepared  by  the  gradual  addition  of  8.2  parts  of  iodine  to  i  part  of  phosphorus 
dissolved  in  carbon  disulphide.  On  gently  distilling  off  the  carbon  disulphide, 
the  iodide  is  left  as  a  yellow  crystalline  solid.  The  compound  melts  at  110°. 

Phosphorus  Tri-iodide,  PI3,  is  obtained  by  employing  a  larger  proportion  of 
iodine  in  the  above  reaction.  It  is  a  solid  substance,  crystallising  in  red  six- 
sided  prisms,  which  are  decomposed  by  water  into  hydriodic  and  phosphorus 
acids. 


OXY  AND  THIO  COMPOUNDS  OF  PHOSPHORUS 
AND  THE  HALOGENS. 

The  following  compounds  are  known,  containing  phosphorus 
combined  with  the  halogens,  and  either  oxygen  or  sulphur — 

POF3;   POC13;  P2O3C14  ;  POBrCl2  j  POBr3. 
PSF3;    PSC13;         —  PSBr3. 

These  compounds  may  be  regarded  as  derived  from  the  haloid 
compounds,  by  the  replacement  of  two  atoms  of  the  halogen  by  an 
equivalent  of  oxygen  or  of  divalent  sulphur  ;  or  they  may  be  viewed 
as  derivatives  of  phosphoric  acid,  by  the  substitution  of  halogen 
elements  in  the  place  of  (HO)  groups.  The  tribasic  phosphoric 
acid,  PO(HO)3,  may  be  regarded  as  giving  rise  to  the  compounds 
POF3,  POC13,  &c.  ;  while  the  compound  P2O3C14  may  be  viewed 
as  a  derivative  of  pyrophosphoric  acid,  P2O3(HO)4. 


468  Inorganic  Chemistry 

Phosphoryl  Fluoride  (phosphorus  oxyfluoride),  POF3,  may  be  obtained  by 
the  action  of  phosphoryl  chloride  (POC13)  upon  zinc  fluoride  — 

3ZnF2  +  2POC13  =  2POF3  +  3ZnCl2. 

It  may  also  be  prepared  by  gently  heating  a  mixture  of  finely  powdered 
cryolite  and  phosphorus  pentoxide— 

2(AlF3,3NaF)  +  2P2O5=4POFa-f  Al2O3  +  3Na2O. 

Phosphoryl  fluoride  is  a  colourless  gas,  which  fumes  in  the  air,  and  is  de- 
composed by  water.  The  gas  in  a  dry  condition  does  not  attack  glass. 

Thiophosphoryl  Fluoride,  PSF3,  is  most  readily  prepared  by  gently  heat- 
ing in  a  leaden  tube  a  mixture  of  dry  lead  fluoride  and  phosphorus  penta- 
sulphide  — 

=3PbS  +  2PSF3. 


The  gas  may  be  collected  over  mercury. 

Thiophosphoryl  fluoride  is  a  colourless  gas,  which  spontaneously  inflames 
when  a  jet  of  it  is  allowed  to  escape  into  the  air.  It  burns  with  a  pale  greenish 
non-luminous  flame.  In  pure  oxygen  the  gas  burns  with  a  yellow  and  more 
luminous  flame.  The  gas  is  decomposed  by  heat  into  phosphorus  fluoride, 
phosphorus,  and  sulphur.  When  heated  in  a  glass  vessel,  sulphur  and  phos- 
phorus are  deposited,  and  silicon  tetrafluoride  is  formed  — 


Phosphoryl  Chloride  (phosphorus  oxychloride\  POC13.—  This 
compound  is  formed  by  the  first  action  of  water  upon  phosphorus 
pentachloride  (q.v^).  It  is  also  obtained  when  phosphorus  penta- 
chloride  and  pentoxide  are  heated  together  in  a  sealed  tube  — 


5POC13. 


It  is  most  readily  prepared  by  heating  phosphorus  pentachloride 
with  either  oxalic  acid  or  boric  acid,  thus  — 

PC15+  H2C2O4  =  POC13  +  2HC1  +  CO2  +  CO. 


Properties.  —  Phosphoryl  chloride  is  a  colourless  fuming  liquid. 
which  boils  at  107.23.°  When  cooled  to  about  —  10°  it  solidifies  to 
a  white  crystalline  mass,  which  melts  at  0.8°.  It  is  decomposed  by 
water  with  formation  of  tribasic  phosphoric  acid  and  hydrochloric 
acid  — 

=  PO(HO)3 


Oxides  and  Oxy acids  of  Phosphorus  469 

Pyropliosphoryl  Chloride,  P2O3C14,  is  obtained  by  passing  nitrogen  peroxide 
through  phosphorus  trichloride,  and  subsequently  distilling  the  liquid.  The 
reaction  is  complicated,  and  cannot  be  expressed  by  a  single  equation  ;  nitro- 
gen is  evolved,  and  phosphorus  pentoxide,  nitrosyl  chloride,  and  phosphoryl 
chloride  are  simultaneously  formed.  Pyrophosphoryl  chloride  is  a  colourless 
fuming  liquid,  boiling  between  210°  and  215°.  It  is  decomposed  by  water, 
and  forms  hydrochloric  acid  and  orthophosphoric  acid  (not  pyrophosphoric 
acid}  — 

P203C14 + 5H40 = 2  H3P04  +  4HC1. 

It  is  converted  by  phosphorus  pentachloride  into  phosphoryl  chloride — 
P203C14  +  PC15=3POC13. 

TMophosphoryl  Chloride,  PSC13,  is  prepared  by  heating  a  mixture  of 
phosphorus  pentasulphide  and  pentachloride — 

3PC15+P2S5=5PSC13. 

It  is  a  colourless  liquid,  boiling  at  125°.  It  fumes  in  moist  air,  being  de- 
composed by  water  into  sulphuretted  hydrogen,  phosphoric  and  hydrochloric 
acids — 

PSC13+4H2O=H2S  +  H3PO4+3HC1. 


OXIDES   AND    OXYACIDS    OF    PHOSPHORUS. 

Four  compounds  of  phosphorus  and  oxygen  are  known,  all  of 
which  are  formed  when  phosphorus  is  burned  in  a  limited  supply 
of  air — 

Phosphorus  monoxide      .         .         .         .         .  P^O  ? 

Phosphorous  oxide  (phosphorus  trioxide)        .  P4O6. 

Phosphorus  tetroxide P2^4- 

Phosphoric  oxide  (phosphorus  pentoxide)       .  P2O5. 

The  two  compounds,  phosphorus  trioxide  and  pentoxide,  are  the 
best  known  of  these  oxides,  and  they  give  rise  respectively  to 
phosphorous  and  phosphoric  acids.  The  following  oxyacids  of 
phosphorus  are  known — 

Corresponding 
Oxide. 

Hypophosphorous  acid  H3PO2 

Phosphorous  acid  .         .  H3PO3  or  P(HO)3         .       P4O0 

Orthophosphoric  acid    .  H3PO4  „  PO(HO)3      } 

Pyrophosphoric  acid     .  H4P2O7),  P2O3(HO)4          P2Ofi 

Metaphosphoric  acid     .  HPO3   „  PO2(HO) 


470 


Inorganic  Chemistry 


When  phosphorus  is  dissolved  in  a  solution  of  aqueous  alcoholic  potash,  and 
dilute  hydrochloric  acid  is  added,  a  yellow  or  reddish  precipitate  is  obtained 
which  was  believed  to  have  the  composition  P4O.  Recent  investigations,  how- 
ever, seem  to  prove  that  the  substance  so  obtained  is  identical  with  red  phos- 
phorus. (Chem,  Soc.  Journal,  Nov.  1899,  and  Nov.  1901.) 

PhosphOPOUS  Oxide  (phosphorous  anhydride},  P4O6  ;  molecular 
weight  =  220. — This  oxide  is  obtained,  mixed  with  a  large  excess 
of  the  pentoxide,  when  phosphorus  is  burned  in  a  tube  through 
which  a  regulated  stream  of  air  is  passed.  In  order  to  obtain  the 
compound  in  a  state  of  purity,  the  following  method  is  employed. 
A  quantity  of  phosphorus  is  introduced  into  a  glass  tube  bent  in  the 


FIG.  124. 

manner  indicated  in  Fig.  124,  and  fitted  into  one  end  of  a  long, 
wide,  Liebig's  condenser.  Into  the  end  of  the  condenser  nearest  to 
the  U-tube  there  is  introduced  a  loose  plug  of  glass  wool,  which 
serves  to  arrest  the  pentoxide.  The  phosphorus  is  ignited  at  the 
open  end  of  the  glass  tube,  and  a  stream  of  air  drawn  through 
the  apparatus  by  means  of  an  aspirator.  A  stream  of  water,  at 
60°,  is  circulated  through  the  condenser,  when  the  easily  fusible 
phosphorous  oxide  is  carried  over,  and  condenses  in  the  U-tube, 
which  is  immersed  in  a  freezing-mixture. 

Properties. — Phosphorous  oxide,  as  it  collects  in  the  cooled 
tube,  is  a  snow-white  crystalline  solid,  which  melts  at  22.5°  to  a 
colourless  liquid.  The  liquid  solidifies  at  2 1°  to  a  white,  waxy-looking 
mass,  consisting  of  monoclinic  prismatic  crystals.  The  liquid  boils 
at  173.1.  It  possesses  an  unpleasant  garlic  smell,  and  is  highly 
poisonous.  Phosphorous  oxide  is  only  very  slowly  acted  upon  by 


Phosphoric  Oxide 


471 


cold   water,  which   gradually  dissolves  it,   forming  phosphorous 
acid — 

P4O6  +  6H2O  =  4H3PO3. 

In  contact  with  hot  water  a  violent  action  takes  place,  in  which 
spontaneously  inflammable  phosphoretted  hydrogen  is  evolved,  and 
a  red  deposit,  consisting  of  red  phosphorus,  is  formed. 

W7hen  heated  in  a  sealed  tube  to  a  temperature  of  440°,  phos- 
phorous   oxide  is  decomposed    into  phosphorus 
tetroxide,  and  red  phosphorus — 


When  exposed  to  air  or  oxygen,  phosphorous 
oxide  is  gradually  oxidised  into  phosphorus 
pentoxide,  but  when  placed  in  warm  oxygen  it 
bursts  into  flame.  When  brought  into  chlorine 
it  also  spontaneously  inflames. 

Phosphorus  Tetroxide,  P2O4.— This  substance  is 
obtained  when  phosphorous  oxide  is  heated  in  a  sealed 
tube  to  a  temperature  of  440°.  It  forms  brilliant  trans- 
parent ciystals,  which  appear  as  a  sublimate  in  the  tube. 
This  oxide  is  highly  deliquescent,  and  dissolves  in  water 
with  evolution  of  heat. 


FIG.  125. 


Phosphoric  Oxide,  P2O3  (or  P4O10).— This  oxide  is  the  main 
product  of  the  combustion  of  phosphorus  in  air  or  oxygen.  It  may 
readily  be  obtained  by  burning  a  quantity  of  phosphorus  in  a  small 
capsule,  and  covering  the  whole  with  a  large  bell-jar  (Fig.  125). 
The  white  clouds  of  phosphoric  oxide  (or  phosphorus  pentoxide) 
collect  as  a  soft  snow-like  substance. 

Properties. — Phosphorus  pentoxide  is  a  white,  amorphous,  and 
very  voluminous  powder.  It  is  without  smell,  although  as  usually 
prepared  it  frequently  possesses  a  slight  garlic  odour,  owing  to  the 
presence  of  phosphorous  oxide.  At  a  temperature  short  of  a  red 
heat  this  oxide  vaporises,  and  recent  determinations  of  its  vapour- 
density  point  to  the  conclusion  that  the  compound  under  these 
conditions  has  a  composition  expressed  by  the  formula  P4O10. 

Phosphoric  oxide  is  extremely  hygroscopic,  absorbing  moisture 
from  the  air  with  great  rapidity.  It  must  therefore  be  preserved 
either  in  well-fitting  stopper  bottles  or  in  hermetically  sealed 
vessels.  Its  affinity  for  water  constitutes  it  the  most  useful  de- 
siccating agent  known  to  chemists  :  prolonged  exposure  to  phos- 


472  Inorganic  Chemistry 

phoric   oxide  removes    the   last   traces   of  aqueous   vapour   from 
gases. 

When  thrown  into  water,  phosphoric  oxide  is  dissolved  with  a 
hissing  sound  resembling  the  quenching  of  hot  iron,  and  forms 
metaphosphoric  acid — 

P205  +  H20  =  2HP03, 

which  gradually  passes  into  orthophosphoric  acid — 
HP03  +  H,0  =  H3P04. 

Phosphoric  oxide  reacts  with  a  number  of  substances,  both 
inorganic  and  organic,  removing  oxygen  and  hydrogen  from  them 
in  the  proportion  in  which  these  elements  form  water  ;  thus,  it 
converts  nitric  acid  into  nitrogen  pentoxide — 

2HNO3-H2O  =  N206. 

In  the  same  way  it  withdraws  the  elements  of  water  from  alcohol, 
with  the  evolution  of  ethylene — 

C2H6O-H2O  =  C2H4. 

Hypophosphorous  Aeid,  H3P(X— This  acid  is  prepared  by 
the  action  of  sulphuric  acid  upon  the  barium  salt — 

Ba(H2P02).j  +  H,SO4  =  BaSO4  +  2H3PO2. 

The  solution,  after  the  removal  of  the  barium  sulphate  by  filtra- 
tion, is  gently  heated  until  its  temperature  rises  to  130°,  when  it  will 
be  sufficiently  concentrated  to  deposit  crystals  when  cooled  to  o°. 

The  barium  hypophosphite  is  obtained  by  boiling  phosphorus 
with  a  solution  of  barium  hydroxide — 

3Ba(HO)2  +  8P  +  6H,O  =  2PH3  +  3Ba(H2PO,),. 

Properties. — Hypophosphorous  acid  is  a  white  crystalline  com- 
pound which  melts  at  17.4°.  When  strongly  heated  it  is  converted 
into  orthophosphoric  acid,  with  the  evolution  of  gaseous  hydrogen 
phosphide — 

2H3P02=H3P04+PH3. 

Hypophosphorous  acid  acts  as  a  powerful  reducing  agent,  on 
account  of  the  readiness  with  which  it  absorbs  oxygen  and  is  con- 
verted into  orthophosphoric  acid — 

HI  HO]  H)  HO) 

HyPO  +  O2=HO[PO;  or  HO  hP  +  O.,-HO  [-PO. 
HOj  HOJ  HOJ  HOJ 


Phosphorous  Acid  473 

Thus  if  hypophosphorous  acid  or  the  sodium  salt  in  solution  be 
added  to  a  solution  of  copper  sulphate,  and  the  mixture  gently 
warmed,  the  copper  is  reduced  even  a  stage  further  than  to  the 
metallic  state,  and  a  dark  red-brown  precipitate  of  copper  hydride^ 
Cu2H2,  is  obtained,  thus  — 


+  3HNaSO4+2Cu2H2. 

This  constitutes  a  characteristic  reaction  for  hypophosphites. 
Hypophosphorous  acid  is  a  feeble  monobasic  acid  ;  its  salts  with 
monovalent  metals  being  represented  by  the  formula  MH2PO2. 

It  is  customary  to  express  the  basicity  of  oxyacids  by  the  number  of  (HO) 
groups  that  are  contained  in  the  molecule,  and  as  this  acid  is  monobasic  its 
constitution  would  be  expressed  by  the  formula  POH2(HO).  Many  of  the 
oxyacids  of  phosphorus,  however,  show  a  tendency  to  exhibit  a  lower  degree 
of  basicity  than  is  represented  by  the  number  of  (HO)  groups  they  contain  ; 
thus,  orthophosphoric  acid,  PO(HO)3,  which  is  tribasic,  and  forms  the  salt 
trisodium  phosphate,  PO(NaO)3,  holds  the  third  atom  of  the  metal  so  loosely 
that  even  such  a  feeble  acid  as  carbonic  acid  is  capable  of  expelling  it  — 

PO(NaO)3+C02  +  H2O  =  PO(HO)(NaO)2+HNaC03. 
or  — 

Na3P04  +  CO2+H2O=HNa2P04+HNaC03. 

The  weaker  acid,  phosphorous  acid,  is  also  tribasic,  P(HO)3,  and  forms 
trisodium  phosphite,  P(NaO)3,  or  Na3PO3.  But  this  salt  is  even  decomposed 
by  water,  into  the  disodium  phosphite,  P(HO)(NaO)2,  or  HNa2PO3. 

Hypophosphorous  acid  being  a  still  weaker  acid,  its  acidic  power  is  destroyed 
as  soon  as  one  atom  of  hydrogen  is  replaced  by  a  base,  and  its  constitution  may, 

H) 
in  harmony  with  these  facts,  be  expressed  by  the  formula  PH(HO)o,  or  HO  VP. 

HOj 

Phosphorous  Acid,  H3PO3,  or  P(HO)3.—  As  already  mentioned, 
this  acid  is  formed  when  phosphorous  oxide  is  dissolved  in  cold 
water. 

It  is  most  readily  prepared  by  the  action  of  water  upon  phos- 
phorus trichloride  — 

PC13  +  3H20  =  3HC1  +  P(HO)3. 

The  production  and  decomposition  of  the  phosphorus  trichloride 
may  be  carried  on  simultaneously,  by  passing  a  stream  of  chlorine 
through  phosphorus  which  is  melted  beneath  water.  The  solution 
is  evaporated  until  its  temperature  rises  to  180°,  when  the  liquid 


474  Inorganic  Chemistry 

will  have  become  so  concentrated  that  on  cooling  it  solidifies  to  a 
crystalline  mass. 

Properties.  —  Phosphorous  acid  is  a  white  crystalline  substance 
which  melts  at  70.1°.  When  heated,  it  decomposes  into  ortho- 
phosphoric  acid,  with  evolution  of  hydrogen  phosphide  — 

4H3PO3  =  3H3PO4+PH3. 

Like  hypophosphorous  acid,  this  compound  absorbs  oxygen,  and 
therefore  is  a  powerful  reducing  agent  ;  silver  salts  are  reduced 
to  metallic  silver,  and  mercuric  salts  are  reduced  to  mercurous 
salts.  By  the  absorption  of  oxygen  it  is  converted  into  ortho- 
phosphoric  acid  — 

H3P03  +  0  =  H3P04. 

Although  a  tribasic  acid,  its  tribasic  salts  are  unstable  ;  the 
sodium  compound,  Na3PO3,  which  is  the  most  stable  inorganic 
salt,  is  decomposed  by  water  into  the  dibasic  salt  — 

Na3PO3+H2O  =  HNa2PO3+NaHO. 
NaCn  HO    ^ 

or  NaO  ^P  +  H2O  =  NaO  iP  +  NaHO. 

NaOJ  NaOj 

Orthophosphorie  Acid,  H3PO4,  or  PO(HO)3.—  This  acid  is 
formed  when  phosphorus  pentoxide  is  dissolved  in  boiling  water, 
or  when  the  solution  of  the  oxide  in  cold  water  is  boiled  — 

P2O5+3H2O  =  2H3PO4. 

Orthophosphorie  acid  is  readily  obtained  by  the  oxidation  of  red 
phosphorus  with  nitric  acid.  Copious  red  fumes  are  evolved,  and 
phosphoric  acid  remains  in  solution. 

Phosphoric  acid  is  prepared  on  a  large  scale  by  the  action  of 
sulphuric  acid  upon  bone  ash,  as  in  the  process  for  the  manu- 
facture of  phosphorus  — 

+  3H2SO4=3CaSO4 


The  calcium  sulphate  is  removed  by  filtration,  and  the  solution 
evaporated  to  a  syrup.  Prepared  in  this  way  the  acid  usually 
contains  arsenic.  This  is  removed  by  first  reducing  it  to  arsenious 
oxide  by  means  of  sulphur  dioxide,  and  after  boiling  off  the  excess 
of  sulphur  dioxide,  precipitating  the  arsenic  as  sulphide  by  means 
of  sulphuretted  hydrogen. 


Pyrophosphoric  Acid  475 

Properties. — The  solution  obtained  by  these  methods  is  either 
concentrated  in  vacuo  or  heated  to  a  temperature  of  140°,  and 
allowed  to  cool,  when  the  acid  is  obtained  in  the  form  of  trans- 
parent six-sided  prismatic  crystals  belonging  to  the  rhombic 
system.  The  substance  is  deliquescent,  and  melts  at  38.6°. 

Phosphoric  acid  is  trfbasic,  and  forms  three  series  of  salts, 
according  as  one,  two,  or  three  of  its  hydrogen  atoms  are  replaced 
by  an  equivalent  of  metal.  Thus,  with  the  metal  sodium  the  three 
salts  are  known — 

Dihydrogen  sodium  phosphate  ....  H2NaPO4. 
Hydrogen  disodium  phosphate  ....  HNa2PO4. 
Trisodium  phosphate  (normal  sodium  phosphate)  Na3PO4. 

The  hydrogen  may  be  replaced  by  its  equivalent  of  more  than 
one  base.  Thus,  the  well-known  compound,  microcosinic  salt,  is 
hydrogen  sodium  ammonium  phosphate,  HNa(NH4)PO4,4H2O. 
The  salt,  which  is  precipitated  when  magnesium  sulphate  (in  the 
presence  of  ammonium  chloride  and  ammonia)  is  added  to  a  solution 
of  a  phosphate,  consists  of  the  compound  ammonium  magnesium 
phosphate  (NH4)MgPO4,6H2O. 

The  heavy  metals  usually  only  form  normal  phosphates.  Thus, 
on  the  addition  of  silver  nitrate  to  a  solution  of  either  of  the  three 
sodium  salts,  the  same  silver  salt  is  precipitated,  namely,  tri- 
argentic  phosphate. 

Na3PO4     +  3AgNO3  =  Ag3PO4  +  3NaNO3. 
*HNa2PO4+3AgNO3  =  Ag3PO4  +  2NaNO3+HNO3. 
H2NaPO4  +  3AgNO3=Ag3P04+NaNO3   +2HNO3. 

Pyrophosphorie  Acid,  H4P2O7,  or  P2O3(HO)4.— This  acid  is 
derived  from  orthophosphoric  acid  by  the  withdrawal  of  one 
molecule  of  water  from  two  molecules  of  the  acid.  This  change  is 
effected  by  heating  the  ortho  acid  to  213° — 

2H3P04-H20  =  H4P207. 

*  Hydrogen  disodium  phosphate,  although  belonging  to  that  class  of  com- 
pounds commonly  called  acid  salts,  on  account  of  the  fact  that  it  still  retains 
a  portion  of  the  replaceable  hydrogen  of  the  acid,  is  strongly  alkaline  in  its 
action  upon  litmus  ;  silver  nitrate  is  a  neutral  compound,  hence  in  this  reaction, 
by  mixing  an  alkaline  and  a  neutral  liquid,  an  acid  liquid  is  obtained,  on 
account  of  the  molecule  of  nitric  acid  that  is  set  free. 


476  Inorganic  Chemistry 

The  formation  of  this  acid  from  two  molecules  of  orthophos- 
phoric  acid  will  be  made  clear  by  the  following  formulas — 

HO   HO  HO   HO  HO  HO  HO  HO 

V  = \/  \/    \/ 

O=P  -;O-HH;-O-P  =  O    =     H2o  +  O=P - o -  p=o. 

Pyrophosphates  are  formed  when  monohydrogen  orthophos- 
phates  are  heated.  Thus,  by  heating  hydrogen  disodium  ortho- 
phosphate,  sodium  pyrophosphate  is  formed — 

2HNa2P04-  H20  -  Na4P2O7. 

When  ammonium  magnesium  phosphate  (see  above)  is  heated 
in  the  same  way  it  loses  water  and  ammonia,  and  is  transformed 
into  magnesium  pyrophosphate,  thus — 

2(NH4)MgP04=Mg2P207  +  H20  +  2NH3. 

Properties. — Pyrophosphoric  acid  is  an  opaque  white  crystal- 
line solid,  readily  soluble  in  water.  Its  aqueous  solution  passes 
slowly  into  orthophosphoric  acid,  the  change  taking  place  rapidly 
on  boiling  ;  a  solution  of  this  acid,  therefore,  cannot  be  concen- 
trated by  boiling. 

The  pyrophosphates  are  stable  salts,  and  their  solutions  may  be 
boiled  without  change  ;  by  boiling  with  acids,  however,  they  are 
converted  into  orthophosphates. 

MetaphosphOPie  Acid,  HPO3  or  PO2(HO).— This  acid  is  formed 
when  phosphorus  pentoxide  is  allowed  to  deliquesce.  It  may  be 
obtained  by  the  abstraction  of  one  molecule  of  water  from  one 
molecule  of  orthophosphoric  acid,  which  is  brought  about  by  heat- 
ing the  tribasic  acid  to  redness — 

H3PO4-H2O  =  HPO3. 

It  is  also  obtained  by  strongly  heating  pyrophosphoric  acid — 
H4P207-H20  =  2HPO3. 

The  sodium  salt  is  obtained  by  strongly  igniting  either  dihydrogen 
sodium  phosphate,  H2NaPO4,  or  hydrogen  sodium  ammonium 
phosphate  (inicrocosmic  salt),  HNa(NH4)PO4;  or  dihydrogen 
sodium  pyrophosphate,  H2Na2P2O7. 

Properties. — Metaphosphoric  acid  is  a  transparent  vitreous 
solid  (frequently  termed  glacial  phosphoric  acid}.  It  is  readily 


Metaphosphoric  Acid 


477 


fusible,  and  is  usually  cast  into  sticks.  At  a  high  temperature  it 
may  be  volatilised.  Metaphosphoric  acid  is  easily  soluble  in 
water,  and  its  solution  is  slowly  transformed  into  orthophosphoric 
acid  ;  this  change  takes  place  rapidly  on  boiling,  and  the  acid 
passes  directly  into  the  tribasic  acid  without  the  intermediate 
formation  of  pyrophosphoric  acid  — 

=  H3PO4. 


Metaphosphoric  acid  is  monobasic,  but  it  possesses  the  remark- 
able property  of  forming  a  number  of  salts  which  may  be  regarded 
as  derived  from  several  hypothetical  polymeric  varieties  of  the 
acid. 

Monometaphosphoric  acid,  HPO3,  forms  monometaphosphates,  NaPO3. 
Dimetaphosphoric  acid,  (HPO3)2       ,,       dimetaphosphates,  K2P2O6. 
Trimetaphosphoric  acid,  (HPO3)3      ,,       trimetaphosphates,  Na3P3C>9. 
Tetrametaphosphoric  acid,  (HPO3)4,  ,,       tetrametaphosphates,  Pb^O^. 
Hexametaphosphoric  acid,  (HPO3)6,  ,,       hexametaphosphates,  NagP6Ols. 

The  three  compounds,  ortho-,  pyro-,  and  metaphosphoric  acids, 
are  readily  distinguished  from  each  other  by  means  of  silver  nitrate 
and  their  action  upon  albumen  :  —  • 


Reagent. 

Orthophosphoric 
Acid. 

Pyrophosphoric 
Acid. 

Metaphosphoiic 
Acid. 

Silver  nitrate  . 
Albumen    .     . 

Canary  yellow 
precipitate      of 
Ag3P04 

No  action 

White  crystalline 
precipitate      of 
Ag4P207 

No  action 

White  gelatinous 
precipitate      of 
AgP03 

Coagulates 

Orthophosphoric  acid  is  also  distinguished  by  giving  a  yellow 
precipitate  of  ammonium  phospho-molybdate  upon  the  addition  of 
excess  of  a  solution  of  ammonium  molybdate  in  nitric  acid  (see 
Molybdenum). 

Compounds  of  Phosphorus  containing  Nitrogen.— By  the  action  of  am- 
monia upon  phosphorus  pentachloride,  and  upon  phosphoryl  chloride  (POC13), 
a  number  of  nitrogen  derivatives  are  obtained.  Thus,  when  gaseous  ammonia 
is  passed  over  phosphorus  pentachloride,  and  the  solid  mass  so  obtained  is 
heated  in  a  stream  of  an  inert  gas  until  the  ammonium  chloride  is  driven 
off,  a  white  insoluble  powder  remains  having  the  composition  represented  by 
the  formula  PN(NH),  to  which  the  name phospham  has  been  given — 

PC15+  7NH3=5NH4CH-  PN(NH). 


478  Inorganic  Chemistry 

Phosphoryl  Triamide,  PO(NH2)3,  is  obtained  by  the  action  of  gaseous 
ammonia  upon  phosphoryl  chloride  — 


3  +  6NH3=PO(NH2)3+3NH4C1. 

When  heated  out  of  contact  with  air,  phosphoryl  triamide  yields  ammonia 
*xA  phosphoryl  nitridet  thus  — 

PO(NH2)3=2NHS+  PON. 

Pyrophosphamic  Acids.  —  Three  of  these  compounds  are  known,  which 
may  be  regarded  as  pyrophosphoric  acid,  in  which  1,2,  and  3  of  the  (HO) 
groups  have  been  replaced  respectively  by  the  group  (NH2),  thus  — 

Pyrophosphoric  acid       ....  P2O3(HO)4. 

Pyrophosphamic  acid    ....  P2O3(HO)3(NH2). 

Pyrophosphodiamic  acid        .         .         .  P2O3(HO)2(NH2)2. 

Pyrophosphotriamic  acid       .         .         .  P2O3(HO)(NH2)3. 

Compounds  of  Phosphorus  with  Sulphur.—  A  number  of  compounds  of 
phosphorus  and  sulphur  have  been  obtained  by  heating  together  varying  pro- 
portions of  sulphur  and  red  phosphorus.  The  following  compounds  are 
known  :  — 

Analogous  Oxides. 

Phosphorus  monosulphide  .         .         .     P4S      .         .       — 
Phosphorus  sesquisulphide         .         .     P4S3    .         .       — 
Phosphorus  trisulphide       .         .         .     P2S3    .         .     P4O6 
Phosphorus  tetrasulphide  (?)       .         .     P2S4    .         .     P2O4 
Phosphorus  pentasulphide  .         .     P2S5     .         .     P2O5 

Phosphorus  Pentasulphide,  P2S5.  —  This  compound  is  the  best-known 
member  of  the  series.  It  is  prepared  by  gently  heating  red  phosphorus  and 
fragments  of  sulphur,  in  the  proportion  required  by  the  formula,  in  a  flask. 
The  elements  combine  with  energy,  and  on  cooling  a  solid  mass  is  obtained. 
This  solid  material  is  then  distilled  in  a  current  of  carbon  dioxide,  when  the 
pentasulphide  is  obtained  in  the  form  of  yellow  crystals.  The  compound  may 
also  be  obtained  by  dissolving  ordinary  phosphorus  and  sulphur  in  the  proper 
proportions  in  carbon  disulphide  and  heating  the  solution  in  sealed  tubes 
to  210°.  On  allowing  the  solution  to  cool,  yellow  crystals  of  the  pentasul- 
phide are  deposited.  Phosphorus  pentasulphide  is  decomposed  by  water  with 
the  formation  of  orthophosphoric  acid  and  the  evolution  of  sulphuretted 
hydrogen— 

P2S5  +  8H20=2H3PO4 


ARSENIC. 

'  Symbol,  As.     Atomic  weight  =  75.     Molecular  weight  =  300. 
Vapour  density  =  150. 

Occurrence. — Arsenic  is  found  in  the  free  state  in  nature, 
usually  in  the  form  of  small  nodules,  more  rarely  as  distinct  crystals. 
In  combination  with  sulphur  it  constitutes  the  minerals  realgar,  or 


Arsenic  479 

ruby  sulphur  ;  As2S2,  and  orpiment,  As2S3.  In  combination  with 
metals,  as  arsenides,  it  occurs  widely  distributed,  the  commonest 
ores  being  arsenical  iron,  FeAs2  and  Fe4As3  ;  kupfernickel,  NiAs 
and  NiAs2  ;  and  tin  'white  cobalt,  CoAs2.  With  metals  and  sulphur 
it  is  met  with  in  such  minerals  as  arsenical  pyrites,  mispickel, 
or  white  mundic,  FeS2,  FeAs2  ;  cobalt  glance,  CoS2,CoAs2  ;  nickel 
glance,  NiS2,NiAs2.  Arsenic  is  present  in  small  quantities  in  most 
samples  of  iron  pyrites,  hence  it  finds  its  way  into  sulphuric  acid 
manufactured  from  pyrites.  It  also  occurs  in  coal  smoke,  being 
derived  from  the  pyrites  contained  in  coal,  and  hence  is  present  in 
the  atmosphere  :  during  the  prevalence  of  yellow  fogs  the  amount 
of  arsenic  present  is  very  appreciable. 

Modes  of  Formation.  —  On  the  small  scale,  arsenic  is  obtained 
by  heating  a  mixture  of  arsenious  oxide,  As4O6,  with  powdered 
charcoal  — 

As4O6+6C  = 


On  a  larger  scale  it  is  usually  obtained  either  from  native  arsenic 
or  from  arsenical  pyrites  ;  the  latter  substance,  when  heated,  gives 
up  arsenic,  and  ferrous  sulphide  is  left  behind  — 

FeS,,FeAs2  =  2As  +  2FeS. 

The  mineral  is  heated  in  long  narrow  horizontal  earthenware 
retorts,  into  whose  mouths  are  fitted  earthenware  receivers.  The 
arsenic  volatilises  and  condenses  in  these  receivers  as  a  compact 
crystalline  solid.  It  is  purified  by  redistillation. 

Properties.—  Arsenic  which  has  been  resublimed  is  a  brilliant 
steel-grey  metallic-looking  substance,  forming  hexagonal  rhombo- 
hedral  crystals,  having  a  specific  gravity  of  5.62  to  5.96.  It  is  very 
brittle,  and  is  a  good  conductor  of  heat  and  electricity.  Arsenic 
begins  to  volatilise  at  100°,  and  rapidly  vaporises  at  a  dark- 
red  heat,  passing  from  the  solid  to  the  vaporous  states  without 
liquefying.  The  vapour  has  a  yellow  colour  and  an  unpleasant 
garlic  smell.  When  heated  under  pressure  arsenic  melts  at  500°, 
and  on  cooling  solidifies  to  a  compact  crystalline  mass.  When 
arsenic  is  vaporised  in  a  glass  tube,  in  a  current  of  hydrogen,  it 
condenses  along  the  tube  in  three  distinct  conditions  :  that  which 
is  deposited  nearest  to  the  heated  portion  of  the  tube  is  in  the  form 
of  rhombohedral  crystals  ;  that  which  sublimes  a  little  farthei 
along,  and  condenses  at  a  point  where  the  temperature  is  about 
2io°-22o°,  consists  of  a  black  shining  amorphous  deposit  ;  while  at  a 


480  Inorganic  Chemistry 

still  more  distant  and  cooler  portion  of  the  tube  a  grey  crystalline 
sublimate  is  formed.  These  are  regarded  as  allotropic  modifica- 
tions of  arsenic.  The  amorphous  variety  is  also  formed,  when 
arsenic  hydride  is  decomposed  by  being  passed  through  a  heated 
tube  (q.v.).  Amorphous  arsenic  is  unacted  upon  by  air  at  ordi- 
nary temperatures,  and  only  slightly  oxidised  at  80°.  The  grey 
crystalline  variety  is  readily  oxidised  on  exposure  to  air  at  ordinary 
temperatures. 

Amorphous  arsenic,  when  heated  out  of  contact  with  air  to  360°, 
is  converted  into  the  rhombohedra!  variety. 

Arsenic,  like  phosphorus,  forms  tetratomic  molecules,  its  mole- 
cular weight  as  deduced  from  its  vapour-density  being  75  x  4  =  300. 

When  heated  in  oxygen  arsenic  burns  with  a  bright  bluish-white 
flame,  forming  arsenious  oxide,  As4Oc.  It  is  oxidised  by  sulphuric 
acid,  nitric  acid,  and  other  oxidising  agents.  It  combines  readily 
with  chlorine,  and  when  thrown  into  this  gas  in  the  condition 
of  powder  it  spontaneously  inflames,  forming  arsenic  trichloride. 
Thrown  into  bromine  a  fragment  of  arsenic  spontaneously  in- 
flames, and  burns  as  it  floats  about  upon  the  surface  of  the  liquid. 

Arsenic,  in  many  of  its  characteristics,  resembles  the  true  metals  ; 
it  is  one  of  those  elements  lying  on  the  borderland  between  true 
metals  and  non-metals,  to  which  the  name  metalloid'^  applied.  It 
is  capable  of  forming  alloys  with  metals,  and  an  alloy  of  this 
element  with  lead  is  employed  for  the  manufacture  of  shot.  It  is 
found  that  by  the  addition  of  a  small  proportion  of  arsenic  to  lead 
the  melted  metal  is  more  fluid,  and  therefore  more  readily  assumes 
the  spheroidal  form  when  projected  from  the  shot  tower,  and  on 
solidification  the  alloy  is  considerably  harder  than  pure  lead. 


ARSENIC  HYDRIDE  (Arsenurelted  Hydrogen.     Arsine). 
Formula,  As  H3.     Molecular  weigh  1  =  78.03.     Density  = 


Modes  Of  Formation.—  (  I.)  Arsenic  hydride  is  formed  when 
soluble  arsenic  compounds  are  exposed  to  the  action  of  nascent 
hydrogen:  thus,  when  a  solution  of  arsenious  oxide  is  introduced 
into  a  mixture  from  which  hydrogen  is  being  generated,  such  as 
zinc  or  iron  and  dilute  hydrochloric  or  sulphuric  acid,  arsenic 
hydride  is  obtained,  mixed  with  free  hydrogen  — 


Arsenic  Hydride  481 

(2.)  By  the  same  action  of  nascent  hydrogen,  arsenic  hydride 
is  formed  when  a  solution  of  either  arsenious  oxide,  As4O6,  or 
arsenic  oxide,  As.,0^,  is  subjected  to  electrolysis. 

(3.)  Arsenic  hydride  is  also  formed  when  arsenical  compounds 
are  in  contact  with  organic  matter  which  is  undergoing  decom- 
position. During  the  growth  of  certain  moulds  and  fungi  a  small 
quantity  of  hydrogen  is  evolved,  which  by  its  action  upon  the 
arsenic  compound,  gives  rise  to  the  formation  of  arsenic  hydride. 
By  this  action  arsenic  hydride  is  sometimes  formed  in  dwelling- 
houses  where  arsenical  wall-papers  are  employed,  and  where,  from 
dampness  or  other  causes,  mould  develops.* 

(4.)  Pure  arsenic  hydride  is  prepared  by  the  action  of  dilute 
hydrochloric  or  sulphuric  acid  upon  an  alloy  of  arsenic  and  zinc  — 

As2Zn3  +  3H,SO4  =  2AsH3  +  3ZnSO4, 

or  by  the  action  of  either  water  or  dilute  acid  upon  an  alloy  of 
arsenic  and  sodium,  prepared  by  heating  sodium  in  the  impure 
arsenic  hydride  obtained  by  method  No.  i. 

Properties.  —  Arsenic  hydride  is  a  colourless,  offensive-smell- 
ing, and  highly  poisonous  gas.  Under  pressure  it  condenses  to  a 
colourless  liquid,  which  boils  at  -54.8°  and  solidifies  at  -113.5°. 
The  gas  burns  with  a  lilac-coloured  flame,  forming  water  and  white 
fumes  of  arsenious  oxide  — 

4AsH3  +  6O2  =  As4O6  +  6H2O. 

When  the  supply  of  air  to  the  flame  is  limited,  as  when  a  cold 
surface  is  depressed  upon  it,  water  is  formed  and  arsenic  is  deposited 
as  a  shining  black  amorphous  film  — 


Arsenic  hydride  is  readily  decomposed  by  heat  into  its  elements  : 
thus,  when  the  gas  is  passed  through"  a  glass  tube,  which  is 
heated  at  one  point  by  a  Bunsen  flame,  arsenic  in  the  amor- 
phous condition  is  deposited  upon  the  tube  immediately  beyond 
the  heated  spot.  Even  when  greatly  diluted  with  hydrogen  this 
reaction  takes  place,  and  it  therefore  affords  a  delicate  test  for  the 
presence  of  exceedingly  small  quantities  of  arsenic.  This  method 

*  Extensive  experiments  on  this  subject  by  C.  R.  Sanger  (Proc.  American 
Academy]  have  led  him  to  believe  that  volatile  organic  arsenical  compounds 
are  produced  under  these  circumstances.  No  compound  was  isolated  how- 
ever. 

2  H 


482  Inorganic  Chemistry 

for  the  detection  of  arsenical  compounds  is  known  as  Marstts  test, 
and  may  be  carried  out  by  means  of  the  apparatus  seen  in  Fig. 
126.  Hydrogen  is  generated  in  the  two-necked  bottle  from  zinc 
and  dilute  sulphuric  acid  (which  are  themselves  free  from  arsenic), 
and  the  arsenic  in  the  form  of  an  oxygen  or  a  haloid  compound  is 
introduced.*  On  igniting  the  issuing  gas,  and  depressing  a  white 
porcelain  capsule  into  the  flame,  black  stains  of  amorphous  arsenic 
are  produced  ;  and  if  the  tube  be  heated  as  shown  in  the  figure, 
the  arsenic  is  deposited  as  a  black  film.  The  corresponding  anti- 
mony compound,  SbH3  (y.v.)t  gives  rise  to  a  similar  deposit  of 
metallic  antimony,  when  treated  in  the  same  way  ;  but  the  arsenic 
deposit  is  readily  distinguished  by  being  easily  soluble  in  a  solu- 

tion of  calcium  hypochlo- 
rite.  Many  metals,  such 
as  sodium  or  potassium, 
when  heated  in  arsenic 
hydride,  form  alloys  with 
the  arsenic,  and  hydrogen 
is  set  at  liberty  ;  while 
metallic  oxides  when  simi- 
larly treated  form  metallic 
arsenides  and  water. 
Arsenic  hydride  is  slightly  soluble  in  water,  but  the  solution  on 
exposure  to  air  deposits  arsenic. 

When  passed  into  a  solution  of  silver  nitrate,  metallic  silver  is 
precipitated,  and  a  solution  of  arsenious  oxide  (the  hypothetical 
arsenious  acid,  H3AsO3)  is  obtained,  thus  — 


When  the  gas  is  passed  into  copper  sulphate  solution,  cuprous 
arsenide  is  precipitated  — 


Arsenic  hydride  is  decomposed  by  the  halogens  with  energy, 
forming  the  haloid  compound  of  arsenic  and  the  halogen  acid  — 


*  When  minute  traces  of  arsenic  have  to  be  detected,  as  in  food  analyses, 
.the  material  is  introduced  into  the  cathode  compartment  of  a  small  specially 
constructed  electrolytic  cell  (Thorpe)  in  which  pure  dilute  sulphuric  acid  is 
electrolysed. 


Arsenic  Chloride  483 

Solid  Arsenic  Hydride. — When  arsenide  of  potassium  or  sodium  is  acted 
upon  by  water,  a  soft  brown  solid  substance  separates,  which  contains  equal 
atomic  proportions  of  arsenic  and  hydrogen.  Its  molecular  weight  is  unknown  ; 
its  composition  is  therefore  expressed  by  the  formula,  (AsH)n. 


COMPOUNDS  OF  ARSENIC  WITH  THE  HALOGENS. 
The  following  compounds  are  known  — 

AsF3;     AsCl3;     AsBr3  ;     Asl^ 

Two  other  compounds  with  iodine  have  been  described  contain- 
ing the  elements  in  the  proportion  represented  by  the  formulas, 
AsI2  and  As2I5,  the  molecular  weights  of  which  are  unknown. 

Arsenic  Fluoride,  AsF3,  molecular  weight  =  132,  is  formed  when 
sodium  fluoride  is  heated  with  arsenic  chloride  — 


It  is  best  obtained  by  distilling  a  mixture  of  arsenious  oxide, 
powdered  fluor  spar,  and  sulphuric  acid  in  a  leaden  retort.  The 
hydrofluoric  acid  generated  by  the  action  of  the  acid  upon  the 
calcium  fluoride  reacts  upon  the  arsenious  oxide,  thus  — 

As4O6  +  12H  F  =  4AsF3  +  6H2O. 

Properties.  —  Arsenic  fluoride  is  a  colourless  fuming  liquid, 
boiling  at  60.4°.  It  is  rapidly  decomposed  by  water  into  arsenious 
oxide  and  hydrofluoric  acid.  On  this  account  it  forms  painful 
wounds  when  brought  into  contact  with  the  skin. 

Arsenic  Chloride,  AsCl3,  molecular  weight=  181.35,  is  ob- 
tained when  arsenic  burns  in  chlorine,  or  when  chlorine  is  passed 
over  fragments  of  arsenic  in  a  tube. 

It  is  also  produced  when  either  arsenic  or  arsenious  sulphide  is 
distilled  with  mercuric  chloride  — 

2As  +  6HgCl2  =  3Hg2Cl2  +  2  AsCl3. 
As2S3  +  3HgCl2  =  3HgS  "+2AsCl3. 


It  is  readily  prepared  by  the  action  of  hydrochloric  acid  upon 
arsenious  oxide  ;  for  which  purpose  sodium  chloride,  arsenious 
oxide,  and  sulphuric  acid  are  gently  heated  together  in  a  retort 
connected  with  a  well-cooled  receiver  — 


484  Inorganic  Chemistry 

Properties. — Arsenic  chloride  is  a  colourless,  fuming,  and  some- 
what oily  liquid  which  boils  at  130.2°,  and  is  extremely  poisonous. 
In  the  presence  of  excess  of  water,  or  when  added  to  warm  water, 
it  is  decomposed  into  arsenious  oxide  and  hydrochloric  acid. 
With  a  small  quantity  of  water  a  solid  crystalline  arsenic  chlor 
hydroxide  is  formed,  As(KO)2Cl — 

AsCl3  +  2H20  =  2HCl  +  As(HO)2Cl. 

Arsenious  Bromide,  AsBr3. — This  compound  is  formed  by  the  direct  union 
of  arsenic  with  bromine,  and  is  prepared  by  adding  powdered  arsenic  to  a 
solution  of  bromine  in  carbon  disulphide.  On  evaporation  the  compound  is 
deposited  in  the  form  of  colourless  deliquescent  crystals,  which  melt  at  20°  to 
25°  to  a  straw-coloured  liquid. 

Arsenious  Iodide,  AsI3,  is  obtained  by  heating  a  mixture  of  arsenic  and 
iodine.  It  is  most  conveniently  prepared  by  digesting  a  saturated  ethereal 
solution  of  iodine  with  powdered  arsenic  in  a  flask  with  a  reflux  condenser. 
On  filtering  and  cooling,  the  iodide  deposits  in  the  form  of  lustrous  red  hexa- 
gonal crystals. 


OXIDES   AND    OXYACIDS   OF   ARSENIC. 

Two  oxides  of  arsenic  are  known,  both  of  which  act  as  anhy- 
drides— 

Arsenious  oxide As4O6. 

Arsenic  oxide  (arsenic  pentoxide)  .         .         .     As2O5. 

No  acid  corresponding  to  arsenious  oxide  is  known  in  the  free 
state,  although  the  arsenites  constitute  a  class  of  stable  salts. 

Three  arsenic  acids,  derived  from  arsenic  pentoxide,  are  known, 
analogous  in  constitution  to  the  three  phosphoric  acids,  namely — 

Ortho-arsenic  acid         .         .     H3AsO4  or  AsO(HO)3. 
Pyro-arsenic  acid  .         .         .     H4As2O7  or  As2O3(HO)4. 
Metarsenic  acid    .         .         .     HAs63  or  AsO2(HO). 


ARSENIOUS  OXIDE. 

Formula,  As4O6.     Molecular  weight =396. 

Mode  of  Formation. — Arsenious  oxide  is  formed  when  arsenic 
burns  in  air  or  in  oxygen,  or  when  arsenic  minerals  are  roasted  in 
a  current  of  air.  On  a  small  scale  it  may  be  produced  by  burning 
arsenic  in  a  hard  glass  tube  in  a  stream  of  oxygen,  and  allowing 


Arsenious  Oxide 


485 


che  white  fumes  of  arsenious  oxide  to  pass  into  a  glass  cylinder  (as 
shown  in  Fig.  127),  where  the  greater  part  condenses,  while  the  rest 
is  led  into  a  draught  flue. 

Arsenious  oxide  is  obtained  as  a  secondary  product,  in  the 
metallurgical  process  of  roasting  arsenical  ores  of  nickel,  cobalt, 
tin,  silver,  and  others,  for  the  extraction  of  these  metals.  It 
is  also  obtained  as  a  principal  product  by  roasting  arsenical 
pyrites.  The  ore  is  heated  either  upon  the  hearth  of  a  rever- 
beratory  furnace,  where  it  is  raked  over  from  time  to  time,  or 
it  is  introduced  by  means  of  a  hopper  into  one  end  of  a  long  clay- 
lined  iron  cylinder,  placed  at  an  incline  of  about  I  in  18,  and  caused 
slowly  to  revolve  about  its  longitudinal  axis  (Fig.  128).  The  lower 
end  of  this  cylinder  enters  a  furnace,  the  upper  end  is  connected  to 
a  series  of  brickwork  flues.  The  ore  is  delivered  into  the  upper 
end  of  the  revolving  cylinder,  and  as  it  gradually  gravitates  down 


FIG.  127. 

the  incline,  it  is  completely  roasted  by  the  furnace  flames  which 
pass  over  it,  and  finally  falls  out  into  a  chamber  beneath.  The 
fumes  of  arsenious  oxide  pass  through  a  series  of  chambers  or 
flues,  so  arranged  as  to  present  an  extensive  condensing  surface  to 
the  gases,  and  the  crude  product,  known  as  arsenical  soot,  is  from 
time  to  time  collected.  This  is  known  as  Oxland  and  Hocking's 
revolving  calciner. 

Properties. — Arsenious  oxide,  known  familiarly  as  white  arsenic, 
or  simply  arsenic,  is  known  in  three  modifications — 

(i.)  Amorphous. 

(2.)  Octahedral  crystals  of  the  cubic  or  regular  system. 

(3.)  Prismatic  crystals  of  the  monosymmetric  system. 


486 


Inorganic  Chemistry 


Amorphous  Arsenious  Oxide  is  a  colourless,  transparent,  vitreous 

substance,  which  is  obtained 
when  the  vapour  of  the  oxide 
is  condensed  at  a  temperature 
only  slightly  below  its  vaporis- 
ing point.  On  exposure  it  gra- 
dually becomes  opaque,  being 
transformed  into  the  regular  octa- 
hedral variety.  This  change 
takes  place  from  the  outside,  and 
lumps  of  opaque  "  white  arsenic," 
when  broken,  often  show  a 
nucleus  of  the  vitreous  modifica- 
tion. Amorphous  arsenious  oxide 
may  be  preserved  unchanged 
in  sealed  tubes.  The  change 
from  the  vitreous  to  the  crystal- 
line form  is  attended  with  evolu- 
tion of  heat,  and  a  diminution  of 
specific  gravity  from  3.738  to 
3.689. 

Amorphous  arsenious  oxide, 
when  heated  to  about  200°,  melts, 
and  at  a  higher  temperature 
vaporises.  It  is  soluble  in  108 
parts  of  cold  water. 

Octahedral  Arsenious  Oxide, — 
The  vitreous  form  passes  spon- 
taneously into  this  variety.  It 
is  obtained  directly,  by  quickly 
cooling  the  vapour  of  arsenious 
oxide,  or  by  crystallisation  from 
the  aqueous  solution  of  either 
form  of  the  oxide.  Arsenious 
oxide  is  also  deposited  in  this 
form  from  solution  in  hydro- 
chloric acid. 

Octahedral  arsenious  oxide  is 
less  soluble  in  water  than  the 
amorphous  variety,  I  part  requir- 
ing 355  parts  of  water  for  its 


Arsenic  Pentoxide  487 

solution.  When  heated,  the  crystals  vaporise  without  fusion,  but 
when  heated  under  pressure  they  melt,  and  are  converted  into  the 
vitreous  form. 

Prismatic  Arsenious  Oxide  is  obtained  by  crystallisation  from  a 
hot  saturated  solution  of  arsenious  oxide  in  potassium  hydroxide. 

Aqueous  solutions  of  arsenious  oxide  possess  a  feeble  acid  re- 
action, probably  due  to  the  formation  of  unstable  arsenious  acid, 
H3AsO3.  The  acid  has  not  been  isolated,  and  on  concentration 
the  solution  deposits  crystals  of  arsenious  oxide. 

Arsenious  oxide  is  a  powerful  poison  :  from  2  to  4  grains  usually 
prove  fatal.  It  is  possible,  however,  by  the  habitual  use  of  it,  to 
so  accustom  the  system  to  this  poison,  that  doses  sufficiently  large 
to  cause  certain  death  to  one  unused  to  it  may  be  taken  with 
apparent  impunity.  The  use  of  arsenic  is  said  to  beautify  the 
complexion,  and  to  improve  the  wind.  The  men  who  are  em- 
ployed upon  arsenic  works  are  constantly  liable  to  swallow  doses 
of  arsenious  oxide  which  would  cause  death  to  one  unaccustomed 
to  the  occupation. 

Arsenites.  —  Three  classes  of  arsenites  are  known,  which  may 
be  regarded  as  being  derived  from  the  three  hypothetical  acids  — 

Silver  ortho-arsenite,  Ag3AsO3. 


green), 

Calcium  pyroarsenite,Ca2As2O5. 
Barium  pyro-arsenite, 


/  Potassium  inetarsenite,  KAsO2. 

Metarsenious  acid,  HAsO2,  or  AsO(HO)       J  Acid'P°tassium  j  KAsO  HAsOjj. 

J    metarsemte,    / 
'  Lead  metarsenite,  Pb(AsO2)2. 

The  pigment  known  as  Schweinfurt  green  is  a  double  metar- 
senite and  acetate  of  copper  — 

3Cu(AsO2)2,Cu(C2H3O2)2. 

All  arsenites,  except  those  of  the  alkali  metals,  are  insoluble  in 
water.  When  heated,  most  arsenites  are  converted  into  arsena/es 
and  arsenic  ;  and  when  heated  with  charcoal  the  whole  of  the 
arsenic  is  reduced. 

Arsenic  Pentoxide,  As2O6.—  This  oxide  is  not  formed  when 
arsenic  burns  in  oxygen.  It  is  obtained  by  the  oxidation  of  ar- 


488  Inorganic  Chemistry 

senious  oxide  by  nitric  acid,  and  subsequently  heating  the  arsenic 
acid  so  produced,  to  a  dark-red  heat  — 

2H3AsO4  =  3H2O  +  As2O5. 

Properties.  —  Arsenic  pentoxide  is  a  white  deliquescent  solid, 
completely  soluble  in  water,  with  the  formation  of  arsenic  acid. 

When  strongly  heated  it  breaks  up  into  arsenious  oxide  and 
oxygen  — 


ARSENIC   ACIDS   AND    ARSENATES. 

When  arsenic  pentoxide  is  dissolved  in  water,  crystals  are  ob- 
tained having  the  composition  2AsO(HO)3,H2O.  At  100°  these 
melt  and  lose  water,  leaving  ortho-arsenic  acid,  H3AsO4.  By  the 
withdrawal  of  water  from  this  acid,  both  pyro-  and  metarsenic  acid 
are  obtained. 

Heated  between  140°  and  180°,  two  molecules  of  the  "ortho" 
acid  lose  one  of  water  — 

2H3AsO4=H.4As2O7  +  H2O. 

And  by  heating  the  pyro-arsenic  acid  so  obtained  to  200°,  another 
molecule  of  water  is  expelled,  with  the  formation  of  metarsenic  acid 
(compare  corresponding  acids  of  phosphorus)  — 

H4As2O7  =  2H  AsO3  +  H2O. 

Pyro-  and  metarsenic  acids  are  both  crystalline  solids,  which 
dissolve  in  water  with  the  evolution  of  heat  and  formation  of  ortho- 
arsenic  acid  ;  aqueous  solutions  of  these  two  acids,  therefore, 
cannot  exist.  In  this  respect  they  differ  from  the  corresponding 
phosphorus  acids,  both  of  which  can  be  obtained  in  aqueous 
solution. 

Each  of  the  three  arsenic  acids  forms  salts,  of  which  the  following 
are  examples  :  — 

Trisodium  ortho-arsenate       .         .        .  Na3AsO4. 

Hydrogen  disodium  ortho-arsenate        .  HNa2AsO4. 

Dihydrogen  sodium  ortho-arsenate        .  H2NaAsO4. 

Ammonium  magnesium  ortho-arsenate.  (NH4)MgAsO4. 

Sodium  pyro-arsenate     ....  Na4As2O7. 

Sodium  metarsenate       ....  NaAsCX. 


Arsenic  Trisulphide  489 

The  salts  of  pyro-  and  metarsenic  acids,  like  the  acids  them- 
selves, only  exist  in  the  solid  state  ;  when  dissolved  in  water  they 
pass  into  the  ortho-compounds. 

The  arsenates  are  isomorphous  with  the  corresponding  phos- 
phates. 


COMPOUNDS    OF   ARSENIC    WITH   SULPHUR. 
Three  sulphides  of  arsenic  are  known,  namely — 

Arsenic  disulphide  (found  native  as  Realgar}  .  As2S2. 
Arsenic  trisulphide  (found  native  as  Orpimenf)  .  As2S3. 
Arsenic  pentasulphide  ......  As2S6. 

Arsenic  Disulphide,  As2S2,  is  formed  when  sulphur  and  arsenic, 
or  arsenic  trisulphide  and  arsenic,  are  heated  together  ;  or  by 
heating  arsenious  oxide  and  sulphur — 

As4O6  +  7S  =  2As2S2  +  3SO2. 

It  is  prepared  on  a  large  scale  by  distilling  a  mixture  of  iron 
pyrites  and  arsenical  pyrites — 

FeS2,FeAs2+2FeS2=As2S2  +  4FeS. 

Properties. — Arsenic  disulphide  is  a  red,  vitreous,  brittle  solid, 
having  a  specific  gravity  of  3.5.  It  is  readily  fusible,  and  sublimes 
unchanged.  Heated  in  air  or  oxygen,  it  burns  with  a  blue  flame, 
forming  arsenious  oxide  and  sulphur  dioxide — 

2As2S2  +  7O2  =  4SO2  +  As4O6. 

Arsenic  disulphide  is  employed  in  pyrotechny.  So-called  Bengal 
/ire  consists  of  a  mixture  of  realgar,  sulphur,  and  nitre. 

Arsenic  Trisulphide,  As2S3,  is  obtained  by  heating  sulphur 
and  arsenic  in  the  proportion  required  by  the  formula,  and  sublim- 
ing the  compound. 

It  may  readily  be  produced  by  passing  sulphuretted  hydrogen 
through  a  solution  of  arsenious  oxide  in  hydrochloric  acid — 

As4O6  +  6H2S  =  2As2S3  +  6H2O. 

Properties. — The  compound,  as  obtained  by  precipitation  with 
sulphuretted  hydrogen,  is  a  pure  canary-yellow  solid,  which  easily 


490  Inorganic  Chemistry 

melts,  and  on  again  cooling  forms  a  brittle  crystalline  mass.  It 
volatilises  and  sublimes  unchanged,  but  when  heated  in  air  or 
oxygen  it  burns  with  formation  of  arsenious  oxide  and  sulphur 
dioxide. 

Arsenic  trisulphide  may  be  regarded  as  a  thio-anhydride,  as  it  gives  rise  to 
a  series  of  salts  known  as  thio-arsenites,  or  sulpharsenites.  Thus,  when  arsenic 
trisulphide  is  brought  into  a  solution  of  a  caustic  alkali,  such  as  potassium 
hydroxide,  the  sulphide  readily  dissolves  with  the  formation  of  an  arsenite  and 
thio-arsenite,  thus  — 

As2S3+4KHO=HK2AsO3+HK2AsS3+H2O. 

Upon  the  addition  of  an  acid,  the  salts  are  decomposed  and  arsenic  tri- 
sulphide reprecipitated  — 

HK2As03+HK2AsS3  +  4HCl=4KCl  +  3H20  +  As2S3. 

Thio-arsenites.  —  These  salts  maybe  looked  upon  as  being  derived  from 
three  hypothetical  thio-arsenious  acids,  corresponding  to  the  oxyacids  — 

Ortho-thio-arsenious  acid,  H3AsS3.       Potassium  ortho-thio-arsenite,  K3AsS3. 

f  Ammonium  pyro  -  thio  -  arsenite, 
Pyro-thio-arsenious  acid,  H4As2S5.  -|  (NH4)4As2S3. 

|^  Lead  pyro-thio-arsenite,  Pb2As2S5. 
Meta-thio-arsenious  acid,  HAsS2.  Potassium  meta-thio-arsenite,  KAsS2. 

Thio-arsenites  of  the  alkali  metals,  the  metals  of  the  alkaline  earths,  and  of 
magnesium,  are  soluble  in  water,  but  decompose  on  boiling.  Their  solutions 
are  decomposed  by  acids,  with  evolution  of  sulphuretted  hydrogen  and  pre- 
cipitation of  arsenic  trisulphide,  thus  — 

2K3AsS3  +  6HC1  =  6KC1  +  3H2S  +  As2S5. 

Arsenic  Pentasulphide,  As2S5.—  This  compound  is  prepared 
by  adding  an  acid  to  a  solution  of  a  thio-arsenate,  thus  — 


Arsenic  pentasulphide  is  a  yellow,  easily  fusible  solid.  It  is 
readily  soluble  in  caustic  alkalies,  forming  an  arsenate  and  a  thio- 
arsenate  — 


Arsenic  pentasulphide,  like  the  trisulphide,  gives  rise  to  a  series  of  salts, 
known  as  thio-arsenates.  These  may  be  regarded  as  being  derived  from  the 
three  hypothetical  thio-arsenic  acids  —  • 

(  Tri  potassium  ortho-thio-arsenate,  K3AsS4. 
Ortho-thio-arsenic  acid,  H3AsS4.  -j  Hydrogen  disodium  ortho  -  thio  -  arsenate, 

I.     HNa-jAsS* 

Pyro-thio-arsenic  acid,  H4As2S7.       Magnesium  pyro-thio-arsenate,  Mg2As2S7. 
Meta-thio-arsenic  acid,  HAsS3.       Ammonium  meta-  thio-arsenate  (NH4)AsS3. 


Antimony  491 


ANTIMONY. 

Symbol,  Sb.     Atomic  weight  =  120.2 

Occurrence.  —  Antimony  in  the  uncombined  state  is  found  in 
small  quantities  in  various  parts  of  the  world,  and  notably  in 
Borneo.  In  combination  with  oxygen,  as  Sb2O3,  it  constitutes 
the  mineral  antimony  bloom,  or  white  antimony;  and  as  Sb2O4  it 
occurs  in  antimony  ochre.  In  combination  with  sulphur,  as  Sb2S3, 
it  occurs  as  the  mineral  stibnite,  or  grey  anti  mony  ore,  which  is  the 
most  important  source  of  the  metal  ;  and  with  both  oxygen  and 
sulphur,  as  Sb2O3,2Sb2S3,  it  constitutes  the  mineral  antimony  blende, 
or  red  antimony. 

It  also  occurs  in  combination  with  sulphur  and  with  metals,  in 
the  form  of  thio-anti  'monites. 

Modes  of  Formation.—  (  i.)  Antimony  is  obtained  from  the 
native  sulphide  by  one  of  the  two  following  methods.  The 
broken-up  ore  is  heated  in  plumbago  crucibles  along  with  scrap 
iron.  As  the  mass  melts,  the  sulphur  combines  with  the  iron, 
forming  a  slag  of  iron  sulphide,  and  the  liberated  antimony  settles 
out  beneath  — 

Sb2S3  +  3Fe  =  2Sb  +  3FeS. 


(2.)  The  crude  sulphide  is  first  liquated,  or  melted  in  such  a 
manner  as  to  separate  the  sulphide  from  the  rocky  matter  associated 
with  it.  The  liquated  sulphide  is  then  mixed  with  about  half  its 
weight  of  charcoal,  in  order  to  prevent  the  mass  from  caking,  and 
carefully  roasted.  During  this  process  the  antimony  sulphide  is 
partially  converted  into  antimony  trioxide  (Sb2O3)2,  which  passes 
into  flues,  and  is  there  condensed,  leaving  a  mixture  containing 
antimony  tetroxide  (Sb2O4),  and  unchanged  sulphide.  Most  of 
the  arsenic  present  is  also  oxidised,  and  volatilises  with  the  anti- 
mony trioxide,  while  sulphur  dioxide  escapes.  The  residue,  con- 
sisting of  the  tetroxide  and  sulphide  (known  as  antimony  ash)  is 
mixed  with  an  additional  quantity  of  charcoal  and  with  sodium 
carbonate,  and  heated  to  redness  in  a  crucible,  when  the  changes 
represented  by  the  following  equations  take  place  — 


(i.) 
By  the  action  of  the  carbon  upon  the  sodium  carbonate,  sodium 


492  Inorganic  Chemistry 

is  liberated,  which  combines  with  the  sulphur  of  the  trisulphide, 
forming  sodium  sulphide  and  metallic  antimony  — 

(2.)  Na2CO 
(3.) 


The  sodium  sulphide  in  its  turn  unites  with  a  further  quantity  of 
antimony  sulphide,  forming  a  double  sulphide  of  sodium  and  anti- 
mony, which,  mixed  with  the  sodium  carbonate  and  charcoal, 
constitutes  the  slag.  The  metal  obtained  by  either  process  is 
subsequently  refined. 

Properties.  —  Antimony  is  a  bright,  highly  crystalline,  and  very 
brittle  metal,  possessing  a  bluish-white  colour,  and  having  a  specific 
gravity  of  6.7  to  6.8.  It  is  unacted  upon  by  air  or  oxygen  at  the 
ordinary  temperature,  but  when  heated  it  burns  brilliantly,  forming 
antimony  trioxide.  The  metal  melts  at  630°  ;  and  when  allowed  to 
solidify,  its  crystalline  character  is  seen  by  the  fern-like  appearance 
of  its  surface.  If  a  quantity  of  the  molten  metal  be  allowed  slowly 
to  cool,  and  when  partially  solidified  the  remaining  liquid  portion 
be  poured  off,  the  interior  of  the  mass  is  found  to  be  lined  with 
well-formed  rhombohedral  crystals,  isomorphous  with  arsenic.  In 
the  act  of  solidification  antimony  expands,  a  property  which  it 
imparts  to  its  alloys,  thus  giving  to  them  the  valuable  quality  of 
taking  very  fine  and  sharp  castings.  The  most  important  of  these 
alloys  are  type  metal  (lead  75,  antimony  20,  tin  5)  ;  stereotype  metal 
(lead  112,  antimony  18,  tin  3)  ;  Britannia  metal  (tin  140,  copper  3, 
antimony  9).  Regarded  as  a  metal,  antimony  is  a  bad  conductor 
of  heat  and  electricity. 

Dilute  sulphuric  and  hydrochloric  acids  are  without  action  upon 
antimony.  The  concentrated  acids  convert  it  into  sulphate  and 
chloride  respectively  — 


2O  +  Sb2(SO4)3. 

Antimony  is  oxidised  by  nitric  acid,  dilute  acid  converting  it  into 
antimony  trioxide  or  a  compound  of  the  oxide  with  nitrogen  pent- 
cxide,  Sb2O3,3N2O5,  while  strong  acid  oxidises  it  chiefly  into  anti- 
mony tetroxide  and  pentoxide. 

Powdered  antimony,  when  thrown  into  chlorine,  takes  fire 
spontaneously  and  forms  antimony  trichloride. 

Amorphous  Antimony.  —  Antimony  is  obtained  in  an  amor- 
phous form  by  the  electrolysis  of  a  solution  of  tartar  emetic  in 
antimony  trichloride. 


Antimony  Hydride  493 

Properties.— Amorphous  antimony  presents  the  appearance  of 
a  smooth  polished  rod  of  graphite,  and  has  a  specific  gravity  of 
5.78.  It  always  contains  a  certain  quantity  of  antimony  trichloride 
(from  4  to  12  per  cent.)  ;  but  whether  this  is  in  chemical  union  or 
merely  mechanically  retained  by  the  metal  is  not  known.  Amor- 
phous antimony  is  very  unstable,  and  readily  passes  into  the 
crystalline  modification  ;  a  slight  blow,  even  a  scratch  with  a 
needle,  causes  it  instantly  to  transform  itself  into  the  stable  form 
with  explosive  violence,  the  temperature  at  the  same  moment 
rising  to  250°,  and  clouds  of  the  vapour  of  antimony  trichloride 
being  evolved. 


ANTIMONY  HYDRIDE  (Antimoniuretted  Hydrogen). 
Symbol,  SbH3. 

Modes  of  Formation. — (i.)  This  compound  is  formed  when  a 
solution  of  an  antimony  compound  is  introduced  into  a  mixture 
generating  hydrogen,  such  as  zinc  and  sulphuric  acid. 

(2.)  Hydrogen  containing  as  much  as  u  per  cent,  of  antimony 
hydride  can  be  obtained  by  the  regulated  action  of  an  alloy  of 
antimony  and  magnesium  upon  dilute  hydrochloric  acid  ;  and  on 
cooling  this  gaseous  mixture  in  liquid  air  the  antimony  hydride 
freezes  out  while  the  hydrogen  passes  on. 

Properties.— Antimony  hydride  melts  at  -88°  and  boils  at 
—  17°.  At  ordinary  temperatures  it  is  a  colourless,  offensive- 
smelling,  and  poisonous  gas,  closely  resembling  the  correspond- 
ing arsenic  compound  in  its  general  behaviour.  It  burns  with 
a  violet-tinted  flame,  forming  water  and  antimony  trioxide — 

4SbH3  +  602  =  6H2O  +  Sb4O6. 

When  the  supply  of  air  is  limited  water  is  formed  and  antimony 
is  deposited  ;  when,  therefore,  a  cold  object  is  depressed  upon  the 
flame  black  stains  of  metallic  antimony  are  obtained.  The  gas  is 
easily  decomposed  by  heat,  and  if  passed  through  a  glass  tube 
heated  at  one  point  a  black  deposit  of  antimony  is  formed  upon 
the  glass.  The  antimony  so  deposited  is  insoluble  in  a  solution  of 
bleaching  powder  (see  Arsenic  Hydride).  Antimony  hydride  is 
decomposed  by  the  halogen  elements,  with  the  formation  of  the 
halogen  hydride,  and  the  halogen  compound  of  antimony — 

3  +  3Cl2=3HCl  +  SbCl3 


494  Inorganic  Chemistry 

Sulphuretted  hydrogen,  under  the  influence  of  sunshine,  converts 
antimony  hydride  into  antimony  trisulphide  — 

2SbH3  +  3H2S  =  Sb2S3  +  6H2. 

When  passed  into  silver  nitrate  solution  the  antimony  is  preci- 
pitated in  combination  with  silver,  in  this  way  differing  from  the 
arsenic  analogue  — 


+  3AgNO3=3HNO 


COMPOUNDS  OF  ANTIMONY  WITH  THE  HALOGENS. 

The  compounds  represented  by  the  following  formulae  are 
known  — 

SbF3  ;  SbCl3  ;  SbBr3  ;  SbI3 
SbF6;  SbCl5. 

Antimony  Trifluoiide,  SbF3,  is  prepared  by  dissolving  the  trioxide  in 
aqueous  hydrofluoric  acid.  From  the  concentrated  solution  it  is  deposited  in 
the  form  of  white  deliquescent  crystals.  It  dissolves  in  water,  and  is  gradually 
converted  into  an  oxyfluoride. 

Antimony  Pentafluoride,  SbF5,  is  obtained  when  hydrated  antimony  pent- 
oxide  is  dissolved  in  aqueous  hydrofluoric  acid.  When  the  solution  is  evapo- 
rated the  compound  remains  as  an  amorphous  gum-like  residue. 

Both  of  these  fluorides  exhibit  a  great  tendency  to  unite  with  alkaline 
fluorides,  forming  double  salts,  such  as  SbF3,2KF  ;  SbF3,2NH4F,  in  the  case 
of  the  trifluoride  ;  and  SbF5,KF  ;  SbF5,2KF,  with  the  pentafluonde. 

Antimony  Trichloride,  SbCl3,  is  formed  when  chlorine  is 
passed  over  metallic  antimony,  or  antimony  trisulphide  — 

2Sb  +  3Cl2  =  2SbCl3. 
2Sb2S3  +  9C12  =  4SbCl3  +  3S2C12. 

It  may  also  be  obtained  by  the  action  of  boiling  hydrochloric 
acid,  containing  a  small  quantity  of  nitric  acid,  upon  either  metallic 
antimony,  antimony  trioxide,  or  trisulphide  — 


Properties.  —  Antimony  trichloride  is  a  colourless,  deliquescent, 
crystalline  substance,  melting  at  73.2°  to  an  oily  liquid,  which 
again  solidifies  to  a  soft  translucent  mass.  It  is  soluble  in  alcohol 
and  in  carbon  disulphide,  and  from  the  latter  may  be  crystallised. 
It  may  be  dissolved  in  a  small  quantity  of  water  unchanged.  Thus, 


Antimony  P  entachloride  495 

if  allowed  to  deliquesce  it  liquefies  in  the  water  it  absorbs,  forming 
a  colourless  solution,  which,  upon  evaporation  over  sulphuric  acid, 
again  deposits  crystals  of  the  trichloride.  The  addition  of  larger 
quantities  of  water  results  in  the  formation  of  oxychlorides  *  — 

(i.)     SbCl3+   H90=  2HCl  +  SbOCl. 
(2.)  4SbCl3  +  5H2O  =  10HCl  +  Sb4O5Cl2. 


Continued  boiling  with  water  removes  the  whole  of  the  chlorine, 
forming  the  trioxide  — 


Antimony  chloride  unites  with  alkaline  chlorides,  forming  double  salts  (see 
Antimony  Fluoride),  such  as  SbCl3,2NH4Cl  ;  SbCl3,3KCl.  With  potassium 
bromide  it  forms  the  compound  SbCl3,3KBr,  which,  strangely  enough,  appears 
to  be  identical  with  the  double  compound  of  antimony  tribromide  with 
potassium  chloride,  SbBr3,3KCl. 

Antimony  Pentaehloride,  SbCl5,  is  obtained  by  passing  excess 
of  dry  chlorine  over  metallic  antimony,  or  antimony  trichloride,  in 
a  retort,  when  antimony  pentachloride  distils  over  in  the  excess  of 
chlorine  — 

SbCl3  +  Cl2=SbCl5. 

Properties.  —  Antimony  pentachloride  is  a  nearly  colourless, 
strongly-fuming  liquid.  It  solidifies,  when  cooled,  to  a  mass  of 
colourless  crystals,  which  remelt  at  —6°.  Under  the  ordinary 
atmospheric  pressure  the  pentachloride  dissociates,  when  heated, 
into  the  trichloride  and  chlorine,  but  under  reduced  pressure  it 
may  be  boiled  and  distilled.  Thus,  under  a  pressure  of  22  mm 
it  boils  at  79°. 

By  the  regulated  action  of  ice-cold  water,  oxychlorides  are 
formed  — 


(i.) 

(2.)  SbOCl3  +  H2O  =  SbO2 

Antimony  pentachloride,  and  also  the  oxychlorides,  are  con- 
verted by  hot  water  into  pyro-antimonic  acid  (analogous  to  pyre- 
arsenic  and  pyro-phosphoric  acids)  — 

2SbCl5  +  7H2O  =  H4Sb2O7  +  10H  Cl. 

2HC1. 


*  The  mixed  product  obtained  by  the  action  of  water  upon  antimony  in- 
chloride  is  known  as  powder  of  Algaroth. 


490  Inorganic  Chemistry 

Sulphuretted  hydrogen  (the  sulphur  analogue  of  water)  acts 
upon  antimony  pentachloride,  forming  antimony  sulphotrichloride, 
corresponding  to  the  oxytrichloride — 

SbCl5  +  H2S  =  SbSCl3  +  2HCl. 

Antimony  tribromide,  SbBr3,  and  antimony  tri-iodide,  SbI3,  are  obtained  by 
adding  powdered  antimony  to  solutions  of  the  halogens  in  carbon  disulphide, 
from  which  liquid  the  compounds  are  crystallised :  the  bromide  as  colourless 
deliquescent  crystals,  and  the  iodide  as  hexagonal  ruby-red  crystals.  Both  of 
these  compounds  are  similarly  acted  upon  by  water,  forming  the  oxybromides 
SbOBr  ;  Sb4O5Br2,  and  the  oxyiodides  SbOI ;  Sb4O5I2. 


OXIDES  AND  OXYACIDS   OF  ANTIMONY. 
Three  oxides  of  antimony  are  known — 

Antimony  trioxide  (antimonious  oxide)  .     (Sb2O3)2  or  Sb4O6. 
Antimony  tetroxide          ....     Sb2O4. 
Antimony  pentoxide         ....     Sb2O5. 

No  acids  are  known  corresponding  to  the  trioxide,  although  a  sodium  salt 
of  the  hypothetical  metantimonious  acid,  HSbO2,  has  been  described,  having 
the  composition  NaSbO2,3H2O. 

Three  acids  are  known  derived  from  antimony  pentoxide  which 
are  analogous  to  the  three  arsenic  and  phosphoric  acids — 

Orthoantimonic  acid          ....     H3SbO4. 

Pyroantimonic  acid H4Sb2Or. 

Metantimonic  acid HSbO3. 

Antimony  Trioxide,  Sb4O6,  may  be  prepared  by  the  addition  of 
hot  water  to  a  solution  of  either  antimony  trichloride  or  antimony 
sulphate,  and  washing  the  precipitated  oxide  with  a  solution  of 
sodium  carbonate  to  remove  the  free  acid — 


Properties. — Antimonious  oxide  is  a  white  powder,  which, 
when  volatilised,  condenses  in  two  distinct  forms,  namely,  pris- 
matic crystals  of  the  trimetric  system  and  regular  octahedra.  The 
former  are  deposited  nearest  to  the  heated  material,  the  latter  in 
more  remote  and  cooler  regions.  (See  Arsenious  Oxide,  with 
which  antimonious  oxide  is  isodimorphous.}  Antimonious  oxide 
is  only  very  slightly  soluble  in  water,  and  the  solution  is  without 


Antimony  Pentoxide  497 

action  upon  litmus.  It  is  insoluble  in  nitric  or  sulphuric  acid,  but 
is  dissolved  by  hydrochloric  acid  with  formation  of  the  trichloride. 
It  is  readily  soluble  in  tartaric  acid,  and  in  a  boiling  solution 
of  hydrogen  potassium  tartrate  (cream  of  tartar),  giving  rise  to 
potassium  antimony  tartrate,  or  tartar  emetic — 

4H  K(C4H4OC)  +  Sb4O6  =  4(SbO)K(C4H4O6)  +  2H2O. 
Antimonious  oxide  burns  in  the  air,  forming  the  tetroxide — 
(Sb2O3)2  +  O2  =  2Sb2O4. 

Antimony  Tetroxide,  Sb2O4,  is  formed  when  the  trioxide  burns 
in  air.  It  may  be  prepared  by  strongly  heating  antimony  pent- 
oxide — 

2Sb2O5=2Sb2O4  +  O2. 

Properties. — Antimony  tetroxide  is  a  white  non-volatile  powder 
which  is  insoluble  in  water.  It  is  decomposed  by  boiling  hydrogen 
potassium  tartrate,  forming  tartar  emetic  and  metantimonic  acid, 
thus— 

H  K(C4H406)  +  Sb204  =  (SbO)K(C4H406)  +  HSbO3. 

Antimony  Pentoxide,  Sb2O5,  is  obtained  by  oxidising  metallic 
antimony  with  nitric  acid,  and  heating  the  antimonic  acid  so 
obtained  to  a  temperature  not  exceeding  275°. 

Properties. — Antimony  pentoxide  is  a  straw-coloured  powder, 
insoluble  in  water.  When  heated  to  300°  it  gives  up  oxygen  and 
is  converted  into  the  tetroxide.  Its  feeble  acidic  character  is  seen 
by  its  formation  of  an  alkaline  metantimonate  when  fused  with  an 
alkaline  carbonate — 

Sb2O5+  Na2CO3  =  CO2  +  2NaSbO3. 

Antimonic  Acids  and  Antimonates. — None  of  the  three  antimonic  acids 
can  be  obtained  by  the  action  of  water  upon  the  oxide.  Pyro-antimonic  acid 
is  formed  when  antimony  pentachloride  is  treated  with  hot  water,  and  the 
precipitate  dried  at  100° — 

2SbCl5+ 7H2O= H4Sb2O7  +  10HC1. 

Pyro-antimonic  acid  readily  passes,  by  loss  of  water,  into  metantimonic 
acid — 

H4Sb207  -  H20=2HSb03. 

Metantimonic  acid  is  also  formed  by  oxidising  metallic  antimony  by  means 
of  nitric  acid — 

2Sb+4HNO3=2HSbO3  +  NO2+3NO  +  H2O, 

2  I 


498  Inorganic  Chemistry 

or  by  the  decomposition  of  an  aqueous  solution  of  a  metantimonate  by  means 
of  nitric  acid— 

KSbO3+  HNO3=  KNO3  +  HSbO3. 

On  allowing  the  precipitated  metantimonic  acid  to  remain  for  a  long  time 
in  contact  with  water  it  is  converted  into  ortho-antimonic  acid,  H3SbO4  —  • 

HSb03  +  H20  =  H3Sb04. 


No  salts  of  ortho-antimonic  acid,  H3SbO4,  are  known  ;   the  antimonates, 
therefore,  belong  to  the  two  acids,  pyro-antimonic  acid  and  metantimonic  acid  — 

Pyro-antimonates.*  Metantimonates.* 

Normal  potassium  pyro-anti-  Potassium  metantimonate,  KSbO3. 

monate,  K4Sb2O7. 
Hydrogen    potassium    pyro-  Barium  metantimonate,  Ba(SbO3)2. 

antimonate,  H2K2Sb2O7. 


COMPOUNDS   OF  ANTIMONY    WITH  SULPHUR. 
Two  sulphides  of  antimony  are  known,  namely  — 

Antimony  trisulphide      .  .      .         .         .     Sb2S3. 
Antimony  pentasulphide          .         .         .     Sb2S5. 

Antimony  Trisulphide,  Sb2S3.  —  This  compound  occurs  native 
as  the  mineral  stibnite^  or  grey  antimony  ore.  It  is  prepared  by 
heating  a  mixture  of  powdered  antimony  and  sulphur  (in  propor- 
tion required  by  the  formula)  beneath  a  layer  of  fused  sodium 
chloride  in  a  crucible.  It  is  also  formed  when  sulphuretted  hydro- 
gen is  passed  through  a  solution  of  antimony  trichloride,  or  a 
solution  of  tartar  emetic  — 


Properties.  —  Antimony  trisulphide  as  it  occurs  native,  and  as 
obtained  by  the  direct  union  of  antimony  and  sulphur,  is  a  grey- 

*  As  only  two  types  of  antimonates  are  known,  and  as  the  salts  of  the  type 
MSbO3  are  the  best  known,  the  name  antimonates  was  formerly  applied  to 
them,  and  the  term  met  antimonates  was  given  to  the  salts  belonging  to  the 
other  class.  It  is  better,  however,  to  adopt  the  same  system  of  nomenclature 
for  the  antimony  compounds  as  that  which  is  in  use  for  the  similarly  constituted 
arsenic  and  phosphorus  compounds  — 

Phosphates.  Arsenates.  Antimonates. 
Ortho        .        .     M3PO4                    M3AsO4 

Pyro          .         .     M4P2O7                   M4As2O7  M4Sb2O7 

Meta         .         .     MPOa                     MAsOg  MSbO3 


Antimony  Pentasulphide  499 

black  crystalline  substance  ;  as  prepared  by  precipitation  with 
sulphuretted  hydrogen,  and  subsequently  drying  at  200°,  it  is  a 
brick-red  amorphous  powder,  which,  when  melted  and  slowly 
cooled,  solidifies  in  the  crystalline  form.  Antimony  sulphide  sub- 
limes unchanged  when  heated  in  an  inert  gas,  but  when  heated  in 
air  sulphur  dioxide  is  evolved,  and  antimonious  oxide  and  tetroxide 
are  formed.  Heated  with  hydrochloric  acid,  it  evolves  sulphuretted 
hydrogen,  and  forms  antimony  trichloride  — 

Sb2S3+6HCl  =  2SbCl3  +  3H2S. 

Antimony  Pentasulphide,  Sb2S5,  is  obtained  when  antimony 
pentachloride  is  mixed  with  water,  and  sulphuretted  hydrogen 
passed  through  the  liquid  — 


or  — 

2SbO 


Properties.  —  Antimony  pentasulphide  is  a  dark,  orange-red 
powder,  which,  on  being  heated,  is  decomposed  into  the  trisulphide 
and  free  sulphur. 

Both  of  these  antimony  sulphides  may  be  regarded  as  thio-anhydrides,  for 
although  no  thio-acids  derived  from  them  are  known,  salts  have  been  produced 
which  may  be  viewed  as  derivatives  of  hypothetical  thio-acids.  When  the 
trisulphide  is  either  fused  with  caustic  potash,  or  boiled  in  an  aqueous  solution, 
potassium  thio-antimonite  is  formed  — 

2Sb2S3+4KHO=3KSbS2+KSbO2+2H2O. 

Similarly,  when  antimony  pentasulphide  is  dissolved  in  potassium  hydroxide, 
a  mixture  of  antimonate  and  thio-antimonate  is  obtained  — 


4Sb.2S5  +  18KHO=5K3SbS4 

The  following  are  illustrations  of  the  thio-salts  of  antimony  — 
Sulphide.  HtiCal  Salts' 


f  (Ortho)  H3SbS3.         Potassium  thio-antimonite,  K3SbS3. 
<  (Meta)  HSbS2.  Silver  thio-antimonite,  AgSbS2. 

l(Pyro)  H4Sb2S5.          Lead  thio-antimonite,  Pb2Sb2Sg. 

/Potassium  thio-antimonate,  K3SbS4. 

.  ,ortho,HSsbs,    3d^^:ri  w- 

V.  Barium  thio-antimonate,  Ba3(SbS4)2, 
Only  ortho-thio-antimonatey  are  known. 


500  Inorganic  Chemistry 

BISMUTH. 

Symbol,  Bi.     Atomic  weight  =  208 

Occurrence.  —  Bismuth  occurs  most  commonly  in  the  uncom- 
bined  condition.  It  is  met  with  in  combination  with  oxygen,  as 
Bi2O3,  in  bismuth  ochre  j  and  in  combination  with  sulphur,  as 
Bi2S3,  in  bismuth  glance. 

Mode  Of  Formation.  —  Bismuth  is  principally  obtained  from 
the  native  metal,  and  from  ores  with  which  metallic  bismuth  is 
associated.  The  broken-up  ore  is  liquated  by  being  heated  in 
inclined  iron  pipes,  when  the  bismuth  readily  melts  and  drains 
away. 

Pure  bismuth  can  be  prepared  from  the  crude  metal  thus  ob- 
tained, by  first  dissolving  it  in  nitric  acid,  forming  bismuth  nitrate 
Bi(NO3)3,  and  then  precipitating  the  basic  nitrate  by  the  addition 
of  water  — 


=  (BiO)NO3,H2 

The  basic  nitrate  is  next  dried  and  heated  in  a  crucible  with 
charcoal  ;  the  salt  is  first  converted  into  the  trioxide  by  the  action 
of  heat,  and  the  oxide  is  then  reduced  by  the  carbon  — 

2(BiO)NO3,H2O  =  Bi203  +  N,O4  +  O  +  2H2O. 


Properties.  —  Bismuth  is  a  lustrous  white  metal  with  a  faint 
reddish  tinge.  It  melts  at  268.3°.  If  the  molten  metal  be  allowed 
to  cool  until  partially  solidified,  and  the  remaining  liquid  be  then 
poured  ofif,  obtuse  rhombohedral  crystals  (belonging  to  the  hexa- 
gonal system),  closely  approaching  to  the  cube,  are  obtained. 

The  specific  gravity  of  bismuth  is  9.823  ;  it  is  extremely  brittle, 
and  a  poor  conductor  of  electricity.  Bismuth  is  unacted  upon  by 
dry  air  at  ordinary  temperatures  ;  moist  air  tarnishes  its  surface. 
Heated  in  air  or  oxygen  it  burns,  forming  the  trioxide.  It  is  only 
slightly  attacked  by  hydrochloric  acid,  but  is  converted  by  hot 
sulphuric  acid  into  a  basic  sulphate. 

Bismuth  readily  forms  alloys  with  other  metals,  and  imparts  to 
them  the  useful  properties  of  ready  fusibility  and  hardness.  The 
alloys  known  by  the  general  name  of  fusible  metal  contain  bismuth  ; 
thus,  Wood  's  fusible  metal^  which  melts  at  65°,  consists  of  4  parts 
of  bismuth,  2  of  lead,  I  of  tin,  and  i  of  cadmium. 


Bismuth  DicJiloride  501 

COMPOUNDS  OF  BISMUTH  WITH  THE  HALOGENS. 
Compounds  represented  by  the  following  formulae  are  known — 

BiF3  BiCl3  BiBr3  BiI3. 

(BiCl2)2  (BiBr2)2? 

Bismuth  Trichloride,  BiCl3,  may  be  prepared  by  passing  dry 
chlorine  over  powdered  bismuth  gently  heated  in  a  retort.  A 
yellow  liquid  is  first  formed,  after  which  the  stream  of  chlorine  is 
stopped  and  the  liquid  distilled,  when  the  trichloride  sublimes  in 
the  form  of  crystals.  It  may  also  be  obtained  by  distilling  a  mix- 
ture of  powdered  bismuth  and  mercuric  chloride — 

2Bi  +  6HgCl2  =  3Hg2Cl2  +  2BiCl3. 

Properties.— Bismuth  trichloride  is  a  white,  extremely  deli- 
quescent crystalline  compound.  Heated  in  an  atmosphere  of 
chlorine,  it  melts  to  a  yellow  liquid.  It  is  decomposed  by  water 
with  the  precipitation  of  bismuth  oxychloride — 

BiCl3  +  H2O  =  2HCl  +  BiOCl. 

Bismuth  Diehloride  (BiCl2)2,  is  obtained  by  the  prolonged 
heating  of  mercurous  chloride  and  finely  powdered  bismuth  to  230° 
in  a  sealed  tube.  The  mixture  melts,  and  mercury  collects  at 
the  bottom,  and  on  cooling  the  dichloride  solidifies  as  a  black, 
extremely  deliquescent  solid  upon  the  surface  of  the  mercury. 
When  heated  above  300°  the  dichloride  is  resolved  into  the  tri- 
chloride and  metallic  bismuth.  The  molecular  weight  of  the 
compound  is  unknown. 

Bismuth  Tribromide,  BiBr3,  is  prepared  by  gradually  adding  bromine  to 
powdered  bismuth  and  slightly  warming  the  mixture  for  some  time.  The 
bromide  sublimes  in  the  form  of  golden-yellow,  deliquescent  crystals,  which 
are  decomposed  by  water,  forming  oxybromide,  BiOBr. 

Bismuth  Tri-iodide,  BiI3,  is  prepared  by  subliming  a  mixture  of  iodine  and 
bismuth.  The  sublimate  is  afterwards  finely  powdered  and  again  sublimed, 
and  the  product  finally  distilled  in  a  stream  of  carbon  dioxide,  when  it  forms 
dark  grey  crystals  with  a  bright  metallic  lustre.  Boiling  water  decomposes 
the  compound,  with  formation  of  bismuth  oxyiodide,  BiOI. 

COMPOUNDS   OF   BISMUTH    WITH   OXYGEN. 
Four  oxides  of  bismuth  are  known,  namely — 

Bismuth  dioxide  (Hypobismuthous  oxide) .  .  Bi^ 

,.         trioxide  (Bismuthous  oxide)         .  .  Bi2O3. 

„         tetroxide  (Hypobismuthic  oxide)  .  .  Bi2O4. 

„        pentoxide  (Btsmuthic  oxide)        .  .  Bi2O6. 


502  Inorganic  Chemistry 

None  of  these  compounds  is  an  acid-forming  oxide,  although, 
with  the  exception  of  the  first,  they  all  form  hydrated  oxides. 
These  hydrated  oxides  have  no  acidic  properties,  and  no  salts 
have  been  obtained  in  which  the  acidic  or  negative  portion  of  the 
molecule  consists  of  bismuth  and  oxygen.  All  the  four  oxides, 
when  acted  upon  by  acids,  yield  the  same  series  of  salts  in  which 
the  bismuth  fulfils  the  functions  of  a  trivalent  element,  replacing 
three  atoms  of  hydrogen.  In  the  case  of  the  dioxide,  metallic 
bismuth  is  deposited,  thus  — 

3Bi2O2  +  6H2SO4=  2Bi2(SO4)3  +  2Bi  +  6H2O. 
While  with  the  higher  oxides  oxygen  is  evolved  — 
O3  =  2Bi(NO 


Bismuth  trioxide  is  the  most  stable  and  the  most  important  of 
the  oxides  ;  when  heated  in  air,  the  remaining  three  compounds 
are  converted  into  the  trioxide  :  the  dioxide  by  oxidation,  and  the 
tetroxide  and  pentoxide  by  loss  of  oxygen.  The  trioxide  alone  is 
unchanged  on  being  heated  in  air  or  oxygen. 

Bismuth  Dioxide,  Bi2O2.  —  This  oxide  is  prepared  by  adding  a  mixed  solu- 
tion of  bismuth  trichloride  and  stannous  chloride  to  an  excess  of  a  10  per  cent. 
solution  of  caustic  potash,  air  being  excluded  :  potassium  stannate  is  formed, 
and  bismuth  dioxide  is  precipitated  — 


+  10KHO=Bi202+8KCl+K2Sn03  +  5H20. 

Properties.  —  The  precipitated  compound,  after  being  washed  in  dilute 
caustic  potash  and  dried  in  vacuo,  is  obtained  as  a  black  crystalline  powder. 
When  heated  in  air  it  smoulders,  uniting  with  oxygen  "to  form  the  trioxide. 
When  moist  it  oxidises  spontaneously  — 

Bi2O2-|-O=Bi2O.j. 

Bismuth  Trioxide,  Bi2O3,  is  formed  when  the  metal  is  burnt  in 
air  or  oxygen.  It  may  also  be  obtained  by  heating  the  hydrated 
oxides,  the  carbonate,  or  basic  nitrate,  thus  — 

Bi203,H20  =  Bi203  +  H20. 
Bi2O3,CO2  =  Bi2O3  +  CO2. 


Properties.  —  Bismuth    trioxide   is   a    cream-coloured    powder, 
insoluble  in  and  unacted  upon  bv  water,  and  is  the  only  oxide  of 


Bismuth   Tetr  oxide  503 

bismuth  which  is  unchanged  when  heated  in  the  air  or  in  oxygen. 
It  dissolves  in  acids,  forming  salts  of  bismuth  — 

3H2O  +  2Bi(NO3)3. 


With  small  quantities  of  hydrochloric  acid  it  first  forms  bismuth 
oxychloride,  BiOCl,  which  dissolves  in  additional  acid,  yielding 
the  trichloride  — 

Bi2O   +  2HCl  =  HO  +  2BiOCl. 


None  of  these  compounds  is  soluble  in  water  without  the  presence 
of  excess  of  the  acid.  Water  alone  converts  them  into  insoluble 
basic  salts  and  free  acid,  which  in  the  state  of  extreme  dilution 
is  unable  to  exert  any  solvent  action.  Thus,  in  the  case  of  the 
nitrate  when  water  is  added,  this  compound  is  decomposed  into  the 
basic  nitrate  and  free  nitric  acid  — 

Bi(NO3)3  +  2H2O  =  (BiO)NO3,H2O  +  2HNO3. 

Bismuth  trioxide  forms  three  hydrates,  represented  by  the 
formulae  — 

Bi2O3,H2O.  Bi2O3,2H2O.  Bi2O3,3H2O. 

These  hydrates  have  no  acid  properties,  and  are  incapable  of 
combining  with  bases  to  form  salts,  but  themselves  play  the  part 
of  a  base,  uniting  with  acids  to  form  bismuth  salts. 

The  trihydrate  is  obtained  by  pouring  an  acid  solution  of  bis- 
muth nitrate  into  an  excess  of  strong  aqueous  ammonia  — 


Heated  to  100°  it  is  converted  by  loss  of  water  into  the  mono- 
hydrate  — 

Bi2O3,3H2O  =  Bi2O3,H2O  +  2H2O. 

Bismuth  Tetroxide,  Bi2O4,  is  formed  by  the  action  of  potassium 
hypochlorite  upon  the  trioxide,  the  product  being  dried  at  180°  — 

Bi2O3  +  KC1O  =  Bi2O4  +  KC1. 

Properties.  —  Bismuth  tetroxide  is  a  brownish-yellow  powder, 
which  readily  parts  with  an  atom  of  oxygen  and  passes  into  the 
trioxide. 


504  Inorganic  Chemistry 

Bismuth  PentOXide,  Bi2O5,  is  prepared  by  passing  chlorine  into 
a  nearly  boiling  solution  of  caustic  potash  in  which  is  suspended  a 
quantity  of  bismuth  trioxide  — 


Properties.—  Bismuth  pentoxide  is  a  red  powder,  which  is 
readily  deoxidised  into  the  tetroxide  and  trioxide  by  heat.  It  com- 
bines with  water,  forming  the  hydrate  Bi2O5,H2O,  but  with  excess 
of  water  it  is  gradually  deoxidised  into  hydrates  of  the  tetroxide  or 
trioxide. 

Bismuth  pentoxide  is  reduced,  with  evolution  of  oxygen,  by  both 
nitric  and  sulphuric  acids  — 

Bi2O5  +  3H2SO4  =  Bi2(SO4)3  +  3H2O  +  O2. 

With  hydrochloric  acid  it  behaves  in  the  usual  manner  of 
peroxides,  causing  the  evolution  of  chlorine  — 


Bismuth  TrisulpMde,  Bi2S3.—  This  compound  is  the  only  com- 
pound of  bismuth  with  sulphur  that  is  known  with  certainty.  It 
occurs  native  as  the  mineral  bismuth  glance. 

It  is  precipitated  when  sulphuretted  hydrogen  is  passed  into  a 
solution  of  a  bismuth  salt  — 


It  is  also  obtained  by  heating  together  the  requisite  proportions 
of  bismuth  and  sulphur. 

Properties.  —  As  obtained  by  precipitation,  bismuth  sulphide  is 
a  dark  brown,  almost  black  powder;  the  native  sulphide  forms 
steel-grey  lustrous  crystals. 

It  is  decomposed,  when  strongly  heated,  into  its  constituent 
elements.  Bismuth  sulphide  differs  from  the  corresponding  anti- 
mony and  arsenic  compound  in  not  being  dissolved  by  alkaline 
hydrates  or  sulphides. 


CHAPTER    IV 
THE   ELEMENTS   OF   GROUP    I.   (FAMILY   A.) 

THIS  family  comprises  the  following  five  elements,  known  as  the 

alkali  metals — 

Atomic  Weights.  Melting-points. 

Lithium  (Li)     ....        7.00     ....       180° 
Sodium  (Na)     .         ,         .         .      23.00     ....        95-6° 
Potassium  (K)  .         .         .      39.10     .         .         .         .        62.5° 

Rubidium  (Rb)         .          .          .      85.45     ....         38.5° 
Caesium  (Cs)    ....    132.9       ....        26.5° 

The  most  important  and  the  most  abundant  of  these  elements 
are  potassium  and  sodium,  which  also  were  the  first  to  be  dis- 
covered, having  been  isolated  by  Davy  in  the  year  1807.  The 
element  lithium,  although  widely  distributed  in  nature,  is  for  the 
most  part  found  only  in  minute  quantities;  the  element  was  first 
isolated  by  Bunsen  in  the  year  1855.  The  two  remaining  elements 
are  still  rarer  substances,  usually  met  with  in  very  minute  quantities 
accompanying  sodium  and  potassium.  Both  of  these  elements 
were  discovered  by  Bunsen  by  means  of  the  spectroscope — caesium 
in  1860  and  rubidium  in  the  following  year. 

All  these  elements  are  soft,  silvery-white  metals,  which  may  be 
readily  cut  with  a  knife,  and  which  rapidly  tarnish  in  the  air. 
They  all  decompose  water  at  the  ordinary  temperature.  The 
members  of  this  family  exhibit  that  gradation  in  properties  which 
is  met  with  in  all  similar  families.  Thus,  their  melting-points 
gradually  decrease  as  their  atomic  weights  rise,  as  will  be  seen 
from  the  figures  given  above.  Their  chemical  activity  also  steadily 
increases  as  we  pass  from  lithium  to  caesium.  Thus,  in  the  case 


506  Inorganic  Chemistry 

of  their  behaviour  in  contact  with  water  :  potassium,  when  thrown 
upon  cold  water,  decomposes  that  liquid  with  sufficient  energy  to 
cause  the  ignition  of  the  hydrogen  which  is  evolved ;  sodium 
under  the  same  conditions  melts  and  floats  about  upon  the  sur- 
face, but  the  action  is  not  sufficiently  energetic  to  effect  the 
inflammation  of  the  gas,  unless  the  water  be  previously  heated  ; 
while  with  lithium,  even  with  boiling  water,  the  temperature 
produced  by  the  reaction  does  not  rise  to  the  ignition-point  of 
hydrogen.  The  same  is  also  seen  in  the  spontaneous  oxidation 
of  these  elements  when  they  are  exposed  to  the  air.  Thus, 
lithium  when  cut  with  a  knife,  although  it  is  soon  covered  with 
a  film  of  oxide,  nevertheless  retains  its  bright  metallic  surface  for 
some  seconds  ;  sodium  tarnishes  so  much  more  quickly,  that  the 
film  of  oxide  appears  almost  to  follow  the  knife.  When  potassium 
is  cut  the  bright  surface  can  scarcely  be  seen,  so  rapid  is  the 
oxidation,  and  if  left  exposed  a  fragment  of  the  metal  soon  begins 
to  melt  by  the  heat  of  its  own  oxidation,  and  frequently  spon- 
taneously ignites.  With  rubidium  and  caesium  the  oxidation  is 
even  more  rapid,  and  a  fragment  of  these  metals  freely  exposed  to 
the  air  very  rapidly  takes  fire  spontaneously. 

The  electro-positive  character  of  these  elements  gradually  in- 
creases from  lithium  to  caesium,  which  is  the  most  electro-positive 
of  all  the  known  elements. 

The  term  alkali^  applied  to  metals  of  this  family,  was  originally 
used  (before  any  distinction  was  made  between  potash  and  soda)  to 
denote  the  salt  obtained  by  treating  the  ashes  of  plants  with  water. 
Later  on,  in  order  to  distinguish  between  this  substance  and  what 
became  known  as  the  volatile  alkali  (i.e.  ammonium  carbonate), 
it  was  termed  the  fixed  alkali.  The  first  distinction  between 
potash  and  soda  was  based  upon  the  erroneous  belief  that  the 
former  was  entirely  of  vegetable  origin,  while  the  latter  was  only 
to  be  found  in  the  mineral  kingdom  ;  hence  the  names  vegetable 
alkali  and  mineral  alkali  were  used  to  denote  these  two  sub- 
stances, both  of  which  were  regarded  as  elementary  bodies  until 
1807,  when  Davy  showed  that  they  contained  the  two  metals 
potassium  and  sodium. 


The  Alkali  Metals 


507 


The  resemblance  between  the  different  members  of  this  family 
and  between  their  compounds  is  very  close  ;  so  much  so,  that  in  the 
case  of  sodium,  potassium,  rubidium,  and  caesium,  there  are  scarcely 
any  ordinary  chemical  reactions  by  which  they  can  be  distinguished. 
They  are  all  readily  identified,  however,  by  means  of  the  spectro- 
scope. When  a  minute  quantity  of  a  lithium  salt  is  introduced 
upon  a  loop  of  platinum  wire  into  the  non-luminous  Bunsen  flame, 
the  latter  is  tinged  a  brilliant  crimson-red  colour  ;  a  potassium  salt 


FIG.  129. 

similarly  treated  colours  the  flame  a  delicate  lilac,  while  a  sodium 
compound  gives  a  brilliant  daffodil-yellow  colour.  The  colour 
imparted  to  a  flame  by  rubidium  and  caesium  salts  is  indistinguish- 
able by  the  eye  from  that  given  by  potassium  compounds ;  and, 
moreover,  when  any  of  these  are  mixed  with  a  sodium  salt  the 
intense  yellow  emitted  by  the  latter  completely  masks  the  colours 
given  by  the  others.  By  means  of  the  spectroscope,  not  only  are 
the  apparently  similar  colours  given  by  potassium,  rubidium,  and 


508  Inorganic  Chemistry 

caesium  readily  distinguished,  but  the  presence  of  any  or  all  of 
them  is  easily  detected,  even  when  admixed  with  sodium  salts. 
Spectrum  analysis  is  based  upon  the  fact  that  light  of  different 
colours  has  different  degrees  of  refrangibility,  and  therefore  when 
passed  through  a  prism  the  different  coloured  rays  are  bent  out 
of  their  straight  course  at  different  angles.  Ordinary  white  light 
is  composed  of  rays  of  all  degrees  of  refrangibility,  i.e.  rays  of  all 
colours  ;  hence,  when  a  beam  of  such  light  is  passed  through  a 
prism,  the  various  coloured  rays  are  separated  and  become  spread 
out  in  the  order  of  their  refrangibility,  from  the  least  refrangible 
red  at  the  one  extreme  to  the  deep  violet  at  the  other.  This 
familiar  "rainbow"  coloured  band  of  light  is  termed  the  con- 
tinuous spectrum. 

A  simple  form  of  spectroscope  is  seen  in  Fig.  129.  The  light  is 
caused  to  pass  through  a  narrow  slit  at  the  end  of  the  fixed  tube  B, 
known  as  the  collimator  tube.  If  the  prism  P  be  removed  and  the 
telescope  A  be  moved  round  so  as  to  be  in  a  continuous  line  with  B? 
a  magnified  image  of  the  slit  is  seen  by  the  observer.  When  the 
prism  is  replaced,  and  A  is  moved  into  such  a  position  that  the  bent 
rays  fall  upon  its  lens,  the  continuous  spectrum  is  seen,  which  is  an 
infinite  number  of  strips  of  light  (corresponding  to  the  image  of  the 
slit)  of  all  colours  arranged  side  by  side.  If  the  light  to  be 
examined,  instead  of  being  ordinary  white  light,  were  composed  of 
rays  all  of  one  degree  of  infrangibility  (i.e.  monochromatic  light), 
there  would  be  produced  only  a  single  image  of  the  slit,  which 
would  fall  in  that  position  corresponding  to  the  particular  degree 
of  refrangibility  of  the  light.  Such  a  monochromatic  light  is  pro- 
duced when  a  sodium  salt  is  heated  in  a  Bunsen  flame  ;  if,  there- 
fore, a  salt  of  this  metal  be  introduced  upon  a  loop  of  platinum 
wire  into  the  non-luminous  flame  G,  and  the  light,  after  passing 
through  the  prism,  be  observed  through  A,  instead  of  a  continuous 
spectrum,  there  will  be  seen  a  single  image  of  the  slit,  falling  in 
the  brightest  yellow  part  of  the  spectrum.  When  the  sodium  salt 
is  replaced  by  a  lithium  salt,  it  is  seen  that  two  images  of  the  slit 
are  obtained,  one  in  the  red  and  the  other  in  the  yellow  regions  of 
the  spectrum.  The  light  emitted  from  this  element  consists  of  rays 


The  Alkali  Metals  509 

of  two  degrees  of  refrangibility.  We  say,  therefore,  that  the 
spectrum  of  sodium  is  one  yellow  line*  and  that  of  lithium  con- 
sists of  one  red  and  one  yellow  line.  In  order  to  distinguish  the 
positions  of,  for  example,  the  yellow  lithium  line  and  that  given 
by  sodium,  an  image  of  a  graduated  scale,  illuminated  by  the 
candle  flame  F,  is  also  thrown  into  the  telescope  A. 

If  salts  of  sodium  and  lithium  mixed  together  be  introduced  into 
the  flame  G,  then  three  images  of  the  slit  are  seen,  namely,  the 
yellow  line  given  by  the  sodium,  the  yellow  line  of  the  lithium, 
situated  slightly  nearer  the  red,  and  the  lithium  red  line. 

Potassium,  like  lithium,  gives  a  light  of  two  degrees  of  refrangi- 
bility, forming  consequently  two  images  of  the  slit,  one  in  the 
deep  red  and  the  other  in  the  deep  violet ;  if,  therefore,  lithium, 


FIG.  130. 

sodium,  and  potassium  salts  are  mixed,  and  examined  by  the 
spectroscope,  five  lines  are  seen  (Fig.  130),  namely,  two  red  (one 
belonging  to  lithium  and  one  to  potassium),  two  yellow  (one 
belonging  to  lithium  and  one  to  sodium),  and  the  violet  line  of 
potassium. 

When  analysed  in  this  manner,  the  lights  emitted  by  rubidium 
and  caesium  compounds  are  seen  to  be  totally  different  from  each 
other,  and  from  potassium.  The  spectrum  of  rubidium  consists  of 
two  prominent  lines  in  the  violet  (nearer  the  blue  region  than  that 
belonging  to  potassium),  two  brilliant  red  lines  (very  near  the 
potassium  red  line),  and  a  number  of  less  brilliant  lines  in  the 

*  In  reality,  when  examined  by  a  higher  dispersive  power,  the  sodium  line 
is  seen  to  be  a  group  of  lines. 


510  Inorganic  Chemistry 

orange,  yellow,  and  green.  That  of  caesium  consists  of  two  bril- 
liant blue  lines,  two  bright  red  lines  (near  the  lithium  red  line),  and 
a  number  of  less  prominent  lines  in  the  yellow  and  green.  It  will 
be  seen,  therefore,  that  the  three  elements  potassium,  rubidium, 
and  caesium  may  be  at  once  sharply  distinguished  by  this  optical 
method  of  analysis,  although  they  so  closely  resemble  one  another 
in  their  chemical  behaviour,  as  to  render  it  highly  probable  that 
the  separate  existence  of  the  two  latter  would  never  have  been  dis- 
covered by  chemical  methods  alone. 

Indeed,  before  the  discovery  of  caesium  by  Bunsen,  a  rare 
mineral  known  as  Pollux  (now  known  to  contain  caesium)  was 
mistaken  for  a  potassium  mineral.* 

The  element  lithium,  the  member  of  the  family  that  belongs  to 
the  Typical  series^  exhibits  certain  characteristic  differences  from 
the  other  members.  This  is  seen  particularly  in  the  case  of  the 
carbonate  and  phosphate  of  this  element.  Lithium  carbonate  is  so 
little  soluble  in  water,  that  it  is  precipitated  by  the  addition  of 
carbonate  of  either  sodium  or  potassium  to  a  solution  of  a  lithium 
compound.  The  phosphates  of  all  the  other  members  are  readily 
soluble  in  water,  while  lithium  phosphate  is  almost  insoluble,  and 
is  precipitated  from  solutions  of  a  lithium  salt  by  the  phosphates  of 
either  sodium  or  potassium.  In  these  two  compounds,  the  car- 
bonate and  phosphate,  lithium  behaves  more  like  one  of  the  metals 
of  the  alkaline  earths. 

All  the  metals  of  this  family  are  monovalent,  and  replace  each 
other,  atom  for  atom,  in  chemical  compounds. 


POTASSIUM. 

Symbol,  K.     Atomic  weight  =  39.00. 

OeeUPPence. — In  combination  this  element  is  widely  distributed 
in  nature.  It  forms  an  essential  constituent  of  many  of  the  com- 
mon silicates  and  rocks  which  form  the  earth's  crust.  From 
these  rocks,  by  processes  of  disintegration,  the  potassium  com- 

*  The  student  should  consult  special  works-  on  spectrum  analysis. 


Potassium  $11 

pounds  find  their  way  into  the  soil,  from  whence  they  are  absorbed 
by  plants,  which  can  only  flourish  in  a  soil  that  contains  com- 
pounds of  potassium.  Most  of  the  potassium  found  in  plants  is 
present  in  combination  with  organic  acids. 

From  the  vegetable  kingdom,  potash  passes   directly  into   the 
bodies  of  animals.     The  material  known  as  suint,  which  is  the 


FIG.  131. 

oily  perspiration  of  the  sheep,  that  accumulates  in,  and  is  extracted 
from  the  wool,  consists  of  the  potassium  salt  of  an  organic  acid 
(sudoric  acid).  In  the  form  of  chloride  and  sulphate,  potassium 
is  present  in  sea-water  and  many  mineral  springs.  As  nitrate  it  is 
found  as  a  crystallised  efflorescence  upon  the  soil,  notably  In  Peru 
and  Chili,  where  it  is  associated  with  sodium  nitrate.  The  largest 


512  Inorganic  Chemistry 

supplies  of  potassium  compounds  are  met  with  in  the  great  saline 
deposits  of  Stassfurt,  where  the  element  is  found  as  chloride  (KC1) 
in  sylvine,  as  a  double  chloride  of  potassium  and  magnesium 
(KCl?MgCl2,6H2O)  vs\carnallite^  and  as  a  mixed  sulphate  in  kainite 
(K2S04,MgS04,MgCl2,6H20). 

Modes  of  Formation.  —  (i.)  The  method  by  which  Davy  first 
effected  the  isolation  of  potassium  was  by  the  electrolysis  of 
potassium  hydroxide  :  the  method  may  be  illustrated  by  the  ex- 
periment represented  in  Fig.  131.  A  small  quantity  of  potassium 
hydroxide  is  gently  heated  in  a  platinum  capsule,  which  is  con- 
nected to  the  positive  terminal  of  a  powerful  battery.  A  stout 
platinum  wire,  flattened  out  at  one  end,  is  made  the  cathode. 
When  this  is  introduced  into  the  fused  potash,  a  brisk  evolution 
of  gas  takes  place,  and  minute  beads  of  metallic  potassium  make 
their  appearance  in  the  liquid  and  upon  the  negative  electrode, 
some  of  which  ignite  upon  the  surface.  The  decomposition  takes 
place  according  to  the  equation  — 


(2.)  Potassium  may  also  be  obtained  by  allowing  melted  potassium 
hydroxide  to  pass  over  iron  turnings  heated  to  whiteness,  when  the 
magnetic  oxide  of  iron  is  formed  — 


This  is  known  as  Gay-Lussac  and  Thenard's  method. 

(3.)  The  method  devised  by  Brunner,  and  modified  by  Wohler, 
Deville,  and  others,  consisted  in  heating  to  whiteness  an  intimate 
mixture  of  potassium  carbonate  and  carbon.  This  mixture  was 
obtained  by  first  igniting  in  a  covered  iron  pot  crude  tartar  (hydro- 
gen potassium  tartrate,  or  cream  of  tartar),  which  was  thereby 
decomposed  as  indicated  by  the  equation  — 

2HKC4H4O6  =  K2CO3  +  3C  +  5H2O  +  4CO. 

The  charred  mass  was  then  introduced  into  an  iron  retort,  and 
strongly  heated  in  a  furnace,  when  the  potassium  carbonate  was 
reduced  by  the  carbon,  as  follows'  — 


In  this  process  there  was  frequently  formed  variable  quantities 
of  a  highly  explosive  compound,  owing  to  the  union  of  potassium 
with  carbon  monoxide,  believed  to  have  the  composition  K6(CO)6. 


Potassium 


513 


(4.)  Castner's  process  for  the  manufacture  of  potassium  (1886) 
consisted  in  strongly  heating  potassium  hydroxide  with  a  carbide 
of  iron,  having  approximately  the  composition  CFe2. 

The  potassium  hydroxide,  with  the  powdered  carbide  of  iron,  was 
introduced  into  large  egg-shaped  retorts,  one  of  which  is  repre- 
sented in  Fig.  132.  These  retorts  were  placed  upon  hydraulic 
lifts,  so  that  they  could  be  lowered  away  from  their  covers,  to  the 
ground-level,  in  order  to  be  discharged  at  the  end  of  the  distilla- 
tion. The  retorts  were  heated  by  gaseous  fuel,  and  the  metal,  as 


FIG.  132. 

it  distilled,  was  passed  into  long  narrow  cast-iron  condensers,  from 
which  it  dropped  into  iron  pots,  and  was  protected  from  oxidation 
by  mineral  oil.  The  reaction  which  takes  place  may  be  represented 
by  the  equation — 


(5.)  At  the  present  time  potassium  is  obtained  almost  exclusively 
by  a  modernised  form  of  Davy's  original  method,  namely,  by 
the  electrolysis  of  fused  potassium  hydroxide.  The  process 

2K 


514  Inorganic  Chemistry 

is  conducted  precisely  as  described  for  the  manufacture  of 
sodium. 

Properties. — Potassium  is  a  lustrous  white  metal,  which  at 
ordinary  temperatures  is  sufficiently  soft  to  be  moulded  between 
the  fingers  ;  at  o°  it  is  brittle,  and  shows  a  crystalline  fracture. 
The  metal  is  readily  crystallised  by  melting  a  quantity  of  it  in  a 
vacuous  tube,  and  when  it  has  partially  solidified,  pouring  the  still 
liquid  portion  to  the  other  end  of  the  tube.  Potassium  melts  at 
62.5°,  and  when  boiled  gives  an  emerald-green  vapour.  The 
metal  is  rapidly  acted  on  by  ordinary  air,  its  freshly  cut  surface 
becoming  instantly  covered  with  a  film  of  oxide,  which,  by  absorp- 
tion of  atmospheric  moisture  and  carbon  dioxide,  passes  first  into 
the  hydroxide  and  finally  into  the  carbonate.  Potassium  is  there- 
fore usually  preserved  beneath  naphtha,  or  some  other  liquid 
devoid  of  oxygen. 

When  potassium  is  volatilised  in  a  vacuous  tube,  the  thin  film  of 
metal  which  condenses  upon  the  cool  portion  of  the  tube  is  seen  to 
possess  a  riclvviolet-blue  colour,  when  viewed  by  transmitted  light. 
The  density  of  potassium  vapour  is  about  2o(Dewar  and  Scott),  show- 
ing that  in  the  vaporous  condition  the  molecules  are  monatomic. 

Potassium  dissolves  in  liquefied  ammonia,  forming  a  deep  indigo 
solution  (page  276).  When  potassium  is  thrown  upon  water,  that 
liquid  is  decomposed  with  sufficient  energy  to  cause  the  ignition 
of  the  liberated  hydrogen  (page  172).  When  heated  in  carbon 
dioxide,  potassium  takes  fire,  forming  potassium  carbonate  and 
carbon  (page  304).  Heated  in  carbon  monoxide,  it  forms  the  ex- 
plosive compound  already  mentioned.  Potassium  takes  fire 
spontaneously  in  contact  with  the  halogens,  forming  the  haloid 
compounds  of  the  metal.  When  heated  in  hydrogen  to  360°, 
potassium  hydride,  KH,  is  obtained  as  a  white  crystalline  compound, 
which  is  decomposed  by  moisture  with  evolution  of  hydrogen,  and 
takes  fire  spontaneously  in  oxygen. 

Oxides  Of  Potassium.— When  potassium  is  heated  in  ordinary 
air,  it  takes  fire  and  burns,  giving  rise  to  a  mixture  of  the  oxides 
of  the  metal.  Perfectly  dry  air  or  oxygen  is  without  action  upon 
potassium.  The  most  stable  oxide  is  the  peroxide,  K2O4. 

Potassium  Peroxide,  K2O4,  is  formed  when  potassium  is  burnt 
in  oxygen.  It  may  also  be  obtained  by  heating  the  metal  in 
nitrous  oxide.  It  is  a  yellow  powder  which,  when  strongly  heated, 
evolves  oxygen  and  is  converted  into  a  lower  oxide.  When  thrown 
into  water  a  violent  action  takes  place,  oxygen  and  potassium 


Potassium  Fluoride  515 

hydroxide  being  produced  ;    but  by  the  regulated  action  in  the 
cold  hydrogen  peroxide  is  also  formed  — 


Potassium  Oxide,  K2O,  is  obtained  as  a  greyish-white  mass  by 
the  partial  oxidation  of  potassium  and  subsequent  removal  of  the 
excess  of  metal  by  distillation  in  vacuo.* 

By  the  regulated  combustion  of  potassium  in  nitrous  oxide  Holt  and  Sims 
have  obtained  compounds  having  the  composition  K2O2  and 


Potassium  Hydroxide  (caustic  potash\  KHO,  is  prepared  by 
adding  lime  to  a  dilute  boiling  solution  of  potassium  carbonate,  in 
iron  vessels,  when  calcium  carbonate  is  precipitated  and  potassium 
hydroxide  remains  in  solution  — 

K2CO3  +  Ca(HO),=  CaCO3  +  2KHO, 

the  reaction  being  complete  when  the  addition  of  an  acid  to  a 
small  test  sample  of  the  clear  liquor  produces  no  effervescence. 
This  reaction  is  a  reversible  one,  and  if  the  concentration  is  beyond 
a  certain  limit,  the  potassium  hydroxide  reacts  upon  the  calcium 
carbonate,  reforming  potassium  carbonate.  The  liquid  is  therefore 
constantly  maintained  at  a  certain  state  of  dilution  during  the 
reaction,  at  the  completion  of  which  the  mixture  is  allowed  to 
settle,  and  the  clear  solution  is  then  partially  concentrated  in 
iron  vessels,  and  finally  in  silver,  until  on  cooling  the  substance 
solidifies.  It  is  then  usually  cast  into  sticks.  Potassium  hydroxide 
is  now  also  manufactured  by  the  electrolytic  method  (see  Sodium 
hydroxide). 

Caustic  potash  is  a  white  brittle  solid  ;  it  is  extremely  deliques- 
cent, and  dissolves  in  water  with  evolution  of  heat,  forming  a 
highly  caustic  liquid.  The  solid,  as  well  as  the  solution,  readily 
absorbs  carbon  dioxide,  and  is  employed  in  the  laboratory  for  this 
purpose  when  it  is  desired  to  deprive  a  gas  of  the  last  traces  of  any 
admixed  carbon  dioxide.  A  hot  saturated  solution  of  potassium 
hydroxide,  when  cooled,  deposits  crystals  of  a  hydrate  having  the 
composition  KHO,2H2O. 

Potassium  Fluoride,  KF.  —  This  salt  is  prepared  by  neutralising 
aqueous  hydrofluoric  acid  with  potassium  carbonate,  and  evaporat- 
ing the  solution  in  a  platinum  vessel,  when  the  salt  is  obtained  in 
the  form  of  deliquescent  cubical  crystals.  Potassium  fluoride  dis- 
solves in  aqueous  hydrofluoric  acid  with  evolution  of  heat,  forming 
the  acid  fluoride  of  potassium,  HF,KF,  which  is  obtained  as  an 
anhydrous  salt  when  the  solution  is  evaporated  to  dryness  and 

*  Rengade,  Compt.  Rend.,  1906. 


516  Inorganic  Chemistry 

heated  to  110°.  This  salt  is  not  deliquescent.  When  heated  to  a 
dull  red  heat  it  decomposes  into  the  normal  salt  and  hydrofluoric 
acid  (see  p.  350). 

Potassium  Chloride,  KC1. — This  salt  is  found  in  sea-water,  and 
was  at  one  time  obtained  as  a  secondary  product  in  the  manufacture 
of  bromine  from  sea  salt,  and  of  iodine  from  seaweed,  as  well  as  in 
various  other  industrial  processes.  At  the  present  day  it  is  almost 
exclusively  obtained  from  the  enormous  deposits  of  carnallite  at 
Stassfurt.  The  method  by  which  potassium  chloride  is  obtained 
from  this  double  salt,  KCl,MgCl2,6H2O,  is  based  upon  the  fact, 
that  when  dissolved  in  water  the  salt  dissociates  into  its  two 
constituents  ;  and  when  the  solution  is  concentrated,  the  more 
insoluble  potassium  chloride  first  separates  out,  leaving  the  mag- 
nesium chloride  in  solution. 

In  practice,  the  crushed  crude  carnallite  is  treated  with  boiling 
mother-liquors  from  previous  operations,  in  large  tanks  into  which 
steam  can  be  driven.  These  mother-liquors  are  practically  a 
strong  solution  of  magnesium  chloride,  and  it  is  found  that  while 
potassium  chloride  is  readily  soluble  in  this  liquid,  the  sodium 
chloride  and  magnesium  sulphate  which  are  present  in  the  crude 
carnallite  are  only  slightly  dissolved  by  it,  and  are  therefore  left 
behind  in  the  residue. 

The  muddy  liquid  is  allowed  to  settle  for  about  an  hour,  when  it 
is  drawn  off  into  large  iron  crystallising  tanks.  The  salt  which  is 
then  deposited  contains  from  80  to  90  per  cent,  of  potassium  chloride, 
the  remainder  being  mainly  sodium  and  magnesium  chlorides. 

The  mother-liquor  from  these  crystallising  tanks  is  either  used 
again  for  treating  a  fresh  charge  of  mineral,  or  is  further  evaporated, 
when  crystals  of  carnallite  separate  out ;  for  it  is  found  that  when 
the  amount  of  magnesium  chloride  present  is  greater  than  three 
times  the  proportion  of  potassium  chloride  in  the  solution,  the  liquid 
on  crystallising  deposits  the  double  chloride  of  the  two  metals.  The 
impure  potassium  chloride  from  the  crystallising  tanks  is  purified  by 
washing  with  cold  water,  in  which  the  salt  is  only  slightly  soluble, 
and  by  subsequent  recrystallisation.  Potassium  chloride  crystal- 
lises, like  the  chlorides  of  sodium,  rubidium,  and  caesium,  in  cubes. 

Potassium  Chlorate,  KC1O3. — When  chlorine  is  passed  into  a 
solution  of  potassium  hydroxide,  a  mixture  of  potassium  chlorate 
and  chloride  is  obtained,  thus — 


Potassium  Chlorate 


517 


The  two  salts  in  solution  may  be  separated  by  crystallisation, 
the  chlorate  being  much  less  soluble  in  cold  water  than  the 
chloride. 

On  the  manufacturing  scale,  potassium  chlorate  is  obtained  by 
passing  chlorine  into  milk  of  lime,  when  a  mixture  of  calcium 
chlorate  and  chloride  is  formed  — 


=Ca(ClO3)2+5CaCl2+6H2O. 

The  operation    is  conducted   in  cast-iron    cylinders   connected 
in  series,  one  of  which  is  shown  in  section  in  Fig.  133,  furnished 


FIG.  133. 

with  mechanical  stirring  gear,  a,  b.  b.  The  shaft  and  the  pipes 
conveying  the  chlorine  into  and  from  the  vessel  are  connected  to 
it  by  means  of  the  water-sealed  joints,  c,  e,  e.  The  manhole/"  is  a 
short  wide  leaden  pipe,  dipping  a  few  inches  into  the  liquid,  which 
allows  of  the  periodic  withdrawal  of  samples  for  examination. 
Several  reactions  are  involved  in  the  final  formation  of  the  calcium 
chlorate  ;  in  the  first  case  calcium  hypochlorite  is  produced, 
thus  — 

2  +  CaCl2+2H2O. 


The  calcium  hypochlorite  then  passes  into  a  mixture  of  chlorate 
and  chloride  in  accordance  with  the  equation  — 

3Ca(OCl)2  =  Ca(ClO3)2  +  2CaCl2. 
The  second  change  is  brought  about  by  the  operation  of  two 


518  Inorganic  Chemistry 

causes,  namely,  rise  of  temperature  and  the  presence  of  excess  of 
chlorine.  Heat  alone  is  incapable  of  converting  more  than  a 
small  proportion  of  the  hypochlorite  into  chlorate,  for  the  former 
compound  is  at  the  same  time  decomposed  into  calcium  chloride 
and  free  oxygen.  The  excess  of  chlorine  is  believed  to  act,  through 
the  intervention  of  hypochlorous  acid,  HOC1,  merely  as  a  carrier 
of  oxygen,  reducing  two  molecules  of  calcium  hypochlorite  to 
chloride,  and  oxidising  the  third  to  chlorate,  thus  — 


2CaCl2  +  Ca(ClO3)2  +  2C12  +  2H2O. 

The  absorption  of  chlorine  by  the  milk  of  lime  is  attended  with 
evolution  of  heat  ;  care  is  taken  to  prevent  the  temperature  from 
rising  above  about  70°,  otherwise  loss  results  by  the  decomposition 
of  hypochlorite  with  evolution  of  oxygen,  thus  — 

Ca(OCl)2  =  CaCl2  +  O2. 

When  the  formation  of  calcium  chlorate  is  complete,  the  liquid 
is  allowed  to  settle,  and  is  then  run  into  concentrating  pans,  where 
the  requisite  amount  of  potassium  chloride  in  solution  demanded 
by  the  following  equation  is  added  — 

Ca(ClO3)2  +  2KC1  =  CaCl2  +  2KC1O3. 

The  liquid  is  then  concentrated  in  iron  pans  and  allowed  to 
crystallise,  when  the  moderately  soluble  potassium  chlorate  sepa- 
rate's out,  leaving  the  very  soluble  calcium  chloride  in  solution. 
The  chlorate  is  afterwards  purified  by  recrystallisation. 

Potassium  chlorate,  although  only  moderately  soluble  in  water,  is  much 
more  soluble  in  a  strong  solution  of  calcium  chloride,  hence  there  is  always  a 
loss  (usually  about  10  per  cent.  )  of  chlorate  in  this  process.  Pe'chiney's  pro- 
cess for  obviating  this  consists  in  concentrating  the  liquid  obtained  by  the 
chlorination  of  the  lime  to  a  definite  specific  gravity,  and  then  cooling  it  to 
12°,  when  about  78  per  cent,  of  the  calcium  chloride  crystallises  out.  The 
mother-liquor,  containing  all  the  calcium  chlorate  and  only  the  comparatively 
small  proportion  of  calcium  chloride,  is  then  treated  with  potassium  chloride 
as  usual. 

Like  so  many  of  the  older  manufacturing  processes,  this  for 
the  preparation  of  potassium  chlorate  is  now  being  displaced  by 
modern  electrolytic  methods.  A  solution  of  potassium  chloride  is 


Potassium  Per  chlorate  519 

electrolysed  in  an  undivided  cell  ;  the  anode  consisting  of  a  thin 
sheet  of  platinum,  and  the  cathode  being  a  vertical  grid  of  copper 
wire.  The  solution  is  caused  to  flow  continuously  through  the 
electrolytic  cell,  the  rate  of  flow  being  so  regulated  that  the  tem- 
perature of  the  liquid  is  maintained  at  about  50°  C.,  and  that  the 
proportion  of  chlorate  produced  does  not  rise  above  3  per  cent,  in 
the  liquid.  The  dilute  liquor  is  passed  into  suitable  refrigerators, 
where  the  sparingly  soluble  chlorate  crystallises  out.  The  chemical 
action  may  be  regarded  as  taking  place  in  stages  ;  the  chlorine 
liberated  at  the  anode  there  unites  with  oxygen  and  water,  yielding 
hypochlorous  acid  — 

=  2HC1O. 


At  the  same  time  potassium  hydroxide  is  produced  at  the  cathode, 
with  elimination  of  hydrogen.  The  caustic  potash  coming  in  contact 
with  hypochlorous  acid,  or  with  chlorine,  gives  rise  to  potassium 
hypochlorite,  which  reacting  with  hypocnlorous  acid  produces 
potassium  chlorate  — 

KC1O  +  2HC1O  =  KC1O3  +  2HC1. 

Potassium  chlorate  crystallises  in  white  tables,  belonging  to  the 
monosymmetric  system,  which  when  of  large  size  often  exhibit  fine 
iridescent  colours.  100  parts  of  water  at  o°  dissolve  3.3  parts  of 
the  salt  ;  while  at  100°,  59  parts  are  dissolved. 

Potassium  chlorate  is  used  largely  in  the  manufacture  of  matches, 
on  account  of  the  ease  with  which  it  gives  up  its  oxygen  :  thus^if  a 
small  quantity  of  the  finely  powdered  salt  be  carefully  mixed  with 
an  equally  small  amount  of  red  phosphorus,  the  friction  caused  by 
lightly  rubbing  it  with  a  spatula  is  sufficient  to  cause  the  mixture 
to  detonate  violently.  Similarly,  when  powdered  potassium  chlo- 
rate and  sulphur  are  rubbed  together  in  a  mortar,  the  mixture 
explodes  with  violence.  Potassium  chlorate  is  also  largely  em- 
ployed in  pyrotechny,  especially  in  the  production  of  coloured 
effects,  where  a  fiercely  burning  mixture  is  required. 

Potassium  chlorate  melts  between  360°  and  370°,  and  at  a  tem- 
perature about  380°  begins  to  evolve  oxygen. 

Potassium  Perehlorate,  KC1O4.  —  When  the  chlorate  is  heated, 
it  first  melts  and  begins  to  give  off  oxygen  ;  but  it  soon  begins  to 
partially  solidify,  owing  to  the  formation  of  potassium  perchlorate, 
and  the  evolution  of  oxygen  stops  unless  a  stronger  heat  be 


520  Inorganic  Chemistry 

applied.     The  reaction  at  this  stage  may  be   expressed  by   the 
equation  — 

8KC1O3  =  5KC1O4  +  3KC1  +  2O2. 

The  evolution  of  oxygen,  however,  is  not  an  essential  condition  of 
the  formation  of  the  perchlorate.  By  careful  regulation  of  the  tem- 
perature the  following  decomposition  can  be  made  to  take  place  — 
4KC1O3  =  KC1  +  3KC1O4. 

The  perchlorate  is  separated  by  first  treating  the  residue  with 
cold  water,  which  dissolves  the  greater  part  of  the  chloride,  and 
afterwards  with  warm  hydrochloric  acid,  which  decomposes  any 
remaining  chlorate.  The  salt  is  then  purified  by  crystallisation. 

Potassium  perchlorate  is  very  slightly  soluble  in  cold  water,  100- 
parts  of  water  at  o°  dissolving  only  0.7  part  of  the  salt  ;  while  at 
100°,  20  parts  are  dissolved. 

Potassium  Bromide,  KBr,  and  Iodide,  KL—  These  two  salts 
are  obtained  by  similar  methods.  When  bromine  or  iodine  is 
added  to  a  solution  of  potassium  hydroxide,  the  reaction  which 
takes  place  is  exactly  analogous  to  that  in  the  case  of  chlorine 
(see  Potassium  Chlorate,  above)  — 


If  the  solution  so  obtained  be  evaporated  to  dryness,  and  the 
dry  residue  ignited,  the  bromate  (or  iodate)  is  decomposed,  just  as 
potassium  chlorate  is  decomposed  by  heat,  giving  off  its  oxygeny 
and  being  converted  into  bromide  (or  iodide)  — 
KBrO3  =  KBr  +  3O. 

The  residue,  on  being  dissolved  in   water  and   recrystallised> 
yields  pure  potassium  bromide  (or  iodide). 

These  salts  are  manufactured  by  decomposing  ferrous  bromide, 
Fe3Br8  (or  iodide,  Fe3I8),  with  potassium  carbonate,  thus  — 


The  ferrous  bromide  is  obtained  by  adding  bromine  to  moistened 
iron  borings  (see  Manufacture  of  Bromine). 

Potassium  iodide  and  bromide  both  crystallise  in  cubes,  and  are 
both  readily  soluble  in  water.  These  salts  are  chiefly  used  for 
medicinal  and  photographic  purposes. 

Potassium  Sulphate,  K2SO4.—  This  salt  is  present  in  the  Stass- 
furt  deposits  principally  as  kainite,  K2SO4,MgSO4,MgCl2,6H2O,. 
and  as  polyhalite,  K2SO4,MgSO4,2CaSO4,2H2O.  When  kainite 
is  treated  with  small  quantities  of  water,  or  mother-liquors  from 
other  processes,  the  extremely  soluble  magnesium  chloride  is 


Potassium  Carbonate  521 

removed,  leaving  the  potassium  magnesium  sulphate  ;  and  on. 
adding  to  this  the  requisite  amount  of  potassium  chloride,  the 
following  change  takes  place  — 


K2SO4,MgSO4  +  3KCl  =  2K2SO4  +  KCl,MgCl2. 

From  this  solution  the  potassium  sulphate  crystallises  out. 

Potassium  sulphate  is  also  obtained  by  the  action  of  sulphuric 
acid  upon  the  chloride,  by  a  process  corresponding  exactly  to  the 
first  stage  in  the  Leblanc  soda  process  (^."z/.)  — 

2KC1  +  H2SO4  =  K2SO4  +  2HC1. 

Potassium  sulphate  forms  colourless  rhombic  crystals,  contain- 
ing no  water  of  crystallisation,  therein  differing  from  sodium 
sulphate,  which  crystallises  with  ten  molecules  of  water. 

Potassium  sulphate  is  largely  used  for  agricultural  purposes. 

Potassium  Carbonate,  K2CO3.  —  This  salt  was  formerly  obtained 
exclusively  from  the  ashes  of  wood  and  other  land  plants,  and  was 
known  under  the  name  of  pot-ashes.  The  process  is  still  carried 
on  in  parts  of  Canada  and  the  United  States.  The  wood  is  burned 
in  pits,  and  the  ashes  are  collected  and  lixiviated  with  water 
(with  the  addition  of  a  small  quantity  of  lime)  in  wooden  tubs 
with  perforated  false  bottoms.  The  liquid  which  is  drawn  off  is 
evaporated  to  dryness,  and  usually  calcined  to  burn  away  the 
organic  matter.  This  material,  known  as  American  pot-ashes, 
contains  varying  quantities  of  caustic  potash,  on  account  of  the 
previously  added  lime.  The  so-called  American  pearl-ask  is  a 
purer  product,  obtained  by  concentrating  the  liquor  from  the 
lixiviating  tubs  until  the  less  soluble  impurities  crystallise  out, 
and  finally  evaporating  the  mother-liquor,  containing  the  potassium 
carbonate,  to  dryness,  and  calcining  the  residue. 

Potassium  carbonate  is  also  obtained  from  beet-root  molasses, 
an  uncrystallisable  residue  obtained  in  the  manufacture  of  beet 
sugar,  carried  on  chiefly  in  France.  The  syrup  is  fermented  with 
yeast,  whereby  the  sugar  it  contains  is  converted  into  alcohol,  and 
then  distilled.  The  residual  liquid,  known  as  vinasse,  is  evaporated 
to  dryness  ;  and  from  the  black  residue,  termed  "  vinasse  cinder," 
the  potassium  carbonate  is  extracted. 

Potassium  carbonate  is  obtained  also  from  suint,  which,  as 
already  stated,  contains  considerable  quantities  of  potassium  in 
the  form  of  potassium  sudorate.  The  sheep's  wool  is  lixiviated 


5  22  Inorganic  Chemistry 

with  water,  and  the  solution  evaporated  to  dryness.  The  residue 
is  heated  in  iron  retorts,  whereby  the  organic  potassium  salts  are 
converted  into  carbonate,  while,  at  the  same  time,  ammonia  and 
an  illuminating  gas  are  evolved.  The  carbonaceous  residue  is 
extracted  with  water,  and  the  potassium  carbonate  separated  by 
crystallisation. 

Since  the  development  of  the  Stassfurt  potash  supplies,  these 
sources  of  potassium  carbonate  are  rapidly  sinking  into  the  back- 
ground, and  the  bulk  of  this  compound  is  now  manufactured  from 
potassium  sulphate  by  a  process  similar  to  the  Leblanc  soda 
process  (g.v,}. 

Potassium  carbonate  is  not  manufactured  by  a  method  analogous 
to  the  ammonia-soda  process  (Solvay),  on  account  of  the  too  great 
solubility  of  potassium  bicarbonate  (hydrogen  potassium  carbonate). 

Pure  potassium  carbonate  may  be  obtained  by  igniting  cream 
of  tartar  (see  page  512),  and  extracting  with  water  ;  or  by  heating 
hydrogen  potassium  carbonate,  which  gives  up  water  and  carbon 
dioxide,  thus  — 

2HKCO3=K2CO3  + 


Potassium  carbonate  forms  long  prismatic  crystals  belonging  to 
the  monosymmetric  system,  and  containing  three  molecules  of 
water,  K2CO3,3H2O.  The  anhydrous  salt  is  highly  deliquescent, 
and  very  soluble  in  water. 

Hydrogen  Potassium  Carbonate  (bicarbonate  of  potash\ 
HKCO3,  is  produced  by  passing  carbon  dioxide  into  an  aqueous 
solution  of  the  normal  carbonate,  thus  — 

K2CO3+CO2+H2O  =  2HKCO3. 

This  salt  is  much  less  soluble  in  water  than  the  normal  salt,  and  is 
readily  purilied  by  crystallisation. 

Potassium  Nitrate  (nitre,  saltpetre},  KNO3.—  This  salt  has 
been  known  since  very  early  times.  It  occurs  as  an  efflorescence 
upon  the  earth,  as  a  result  of  the  oxidation  of  organic  nitrogenous 
matter  in  the  presence  of  the  potash  in  the  soil,  and  is  found  in 
the  neighbourhood  of  villages,  more  especially  in  hot  climates, 
where  urine  and  other  readily  decomposable  organic  matters  rich 
in  nitrogen  find  their  way  into  the  surface  soil.  It  has  been  shown 
that  the  process  of  nitrification  which  results  in  the  formation  of 
nitre  under  these  circumstances  is  due  to  the  action  of  specific 
organisms,  or  microbes,  and  never  takes  place  in  their  absence. 


Potassium  Nitrate  523 

At  various  times  this  natural  process  has  been  artificially  carried 
on,  by  mixing  manure  and  other  decomposing  refuse  with  porous 
soil,  lime,  and  wood  ashes,  and  exposing  the  mixture  in  heaps 
which  were  moistened  from  time  to  time  with  drainage  from 
manure.  The  saltpetre  earth,  collected  from  the  natural  sources 
or  from  the  artificial  nitre  plantations,  on  lixiviation  with  water,  and 
subsequent  evaporation,  yielded  crystals  of  potassium  nitrate. 

At  the  present  time  potassium  nitrate  is  almost  exclusively  ob- 
tained from  sodium  nitrate  (Chili  saltpetre},  by  treatment  with 
potassium  chloride  derived  from  the  Stassfurt  supplies.  The  requi- 
site quantities  of  the  two  solutions  are  run  into  a  tank,  and  heated 
by  means  of  steam,  when  the  following  double  decomposition  takes 
place  — 


The  greater  part  of  the  sodium  chloride  is  at  once  precipitated, 
and  is  removed  by  canvas  filters.  The  clear  liquid  is  then  allowed 
to  crystallise  in  tanks  furnished  with  stirring  gear,  in  order  to 
cause  the  formation  of  small  crystals,  and  the  nitre-meal  so  ob- 
tained is  purified  by  recrystallisation. 

Potassium  nitrate  crystallises  usually  in  rhombic  prisms,  but  it 
can  also  be  obtained  in  the  form  of  small  rhombohedral  crystals, 
isomorphous  with  sodium  nitrate. 

The  solubility  of  potassium  nitrate  rapidly  increases  with  rise  of 
temperature  (see  Solubility  Curve,  p.  152).  100  parts  of  water  at 
o°  dissolve  13.3  parts  ;  at  50°,  86  parts  ;  and  at  100°,  247  parts. 

Nitre  melts  at  339°,  and  at  a  higher  temperature  loses  oxygen 
and  is  converted  into  potassium  nitrite  ;  on  this  account  it  readily 
oxidises  many  of  the  elements  when  heated  in  contact  with  them. 
Thus,  a  fragment  of  charcoal  or  sulphur  thrown  upon  melted  nitre 
takes  fire  and  burns  with  great  energy  ;  in  the  one  case  with  forma- 
tion of  potassium  carbonate  and  carbon  dioxide,  and  in  the  other 
of  potassium  sulphate  and  sulphur  dioxide  — 


2K2CO3-f3CO2 
2KNO3+2S=  K,SO4  +  SOj+N,. 

Nitre  is  chiefly  used  in  the  manufacture  of  gunpowder  and  in 
pyrotechny. 

Gunpowder  is  a  mixture  of  nitre,  charcoal,  and  sulphur.     The  proportions 
in  which  these  ingredients  are  present  varies,  within  small  limits,  according 


Inorganic  Chemistry 


to  the  special  kind  of  powder,  as  will  be  seen  from  the  following  table  (Abel 
and  Nobel),  giving  analyses  of  various  powders  manufactured  at  Waltham 
Abbey. 


— 

Fine-Grain. 

Rifle 
Fine-Grain. 

Rifle 
Large-Grain. 

Pebble 
Powder. 

Potassium  nitrate 
,,          sulphate 
Sulphur 
Charcoal     . 
Water 

73-55 
0.36 

10.02 

J4-59 
1.48 

75-04 
0.14 

9-93 
14.09 
0.80 

74-95 
0.15 
10.27 
13-52 
i.  ii 

74.67 
0.09 
10.07 
14.22 

o.95 

These  proportions  are  very  close  to  those  which  would  be  demanded  by  the 
equation  — 


which  was  at  one  time  supposed  to  represent  the  change  which  takes  place 
when  gunpowder  is  exploded.  In  reality  the  decomposition  is  much  more 
complex,  and  it  has  been  shown  that  the  solid  products  consist  of  mixtures  of 
the  following  substances  in  varying  proportions,  depending  upon  the  particular 
powder,  and  the  conditions  of  firing  — 


Potassium  nitrate. 
,,          oxide. 

Ammonium  sesquicarbonate. 
Carbon. 
Sulphur. 

Marsh  gas. 

Oxygen. 

Hydrogen. 


Potassium  carbonate. 
, ,          sulphate. 
,,          sulphide. 
,,          thiosulphate. 
,,          thiocyanate. 

While  the  gases  that  are  evolved  consist  of — 
Carbon  dioxide. 
Nitrogen. 
Carbon  monoxide. 
Sulphuretted  hydrogen. 

From  the  combustion  of  one  gramme  of  powder  the  total  weight  of  solids 
ranges  from  0.55  to  0.58  gramme,  and  the  total  weight  of  the  gaseous  products 
from  0.45  to  0.42  gramme. 

Potassium  Cyanide,  KCN.— This  salt  is  manufactured  from 
potassium  ferrocyanide,  which  is  first  obtained  by  heating  in  an 
iron  pot  a  mixture  of  scrap  iron,  crude  potashes,  and  waste  animal 
refuse,  such  as  hoofs,  horns,  hide,  &c.  The  complex  changes  which 
take  place  do  not  at  once  result  in  the  formation  of  the  ferrocyanide, 
as  this  salt  is  unstable  at  high  temperatures,  but  in  the  production 
of  various  compounds  (the  very  stable  salt  potassium  cyanide 
amongst  them)  which,  when  the  mass  is  subsequently  treated  with 
water,  interact,  with  the  formation  of  potassium  ferrocyanide.  The 


Compounds  of  Potassium  with  Sulphur         525 

aqueous  extract  is  allowed  to  crystallise,  and  the  ferrocyanide  is 
obtained  as  large  lemon-yellow  prisms,  with  three  molecules  of 
water.  When  this  compound  is  dried  and  heated  alone  it  decom- 
poses into  potassium  cyanide,  free  nitrogen,  and  a  carbide  of  iron  — 

K4Fe(CN)6  =  4KCN  +  N2  +  FeC2. 

By  heating  the  ferrocyanide  with  potassium  carbonate  a  larger 
yield  of  the  cyanide  is  obtained,  mixed  with  potassium  cyanate  — 


For  many  commercial  uses  for  which  potassium  cyanide  is 
required  the  presence  of  this  cyanate  is  not  detrimental. 

If  potassium  ferrocyanide  be  heated  with  metallic  sodium  the 
whole  of  the  cyanogen  it  contains  is  converted  into  alkali  cyanide 
(Erlenmeyer)  — 

K4Fe(CN)6  +  2Na  =  4KCN  +  2NaCN  +  Fe. 

The  mixed  potassium  and  sodium  cyanide  thus  obtained  is  well 
suited  for  the  technical  processes  for  which  cyanide  is  required. 

Potassium  cyanide  is  a  white  solid  which  is  extremely  soluble  in 
water,  from  which  it  crystallises  in  white  anhydrous  plates.  When 
heated  the  salt  readily  fuses,  but  is  stable  at  very  high  tempera- 
tures, being  capable  of  being  volatilised  without  decomposition.  In 
the  fused  state  it  is  a  powerful  reducing  agent,  taking  up  oxygen 
to  yield  potassium  cyanate,  KCNO. 


COMPOUNDS   OF   POTASSIUM    WITH   SULPHUR. 
Four  sulphides  of  potassium  have  been  obtained,  namely — 

Potassium  monosulphide  ......  K2S 

Potassium  trisulphide K2S3 

Potassium  tetrasulphide K2S4 

Potassium  pentasulphide  ......  K2S5 

Just  as  potassium  decomposes  water  with  evolution  of  hydrogen 
and  formation  of  potassium  hydroxide,  so  also,  when  heated  in 
sulphuretted  hydrogen  (the  sulphur  analogue  of  water)  it  forms 
potassium  hydrosulphide  (the  analogue  of  potassium  hydroxide) 
and  liberates  hydrogen,  thus — 


526  Inorganic  Chemistry 

When  potassium  hydroxide  and  hydrosulphide  are  mixed  in  equi- 
molecular  proportions,  potassium  monosulphide  and  water  are 
formed  — 


The  liquid,  on  evaporation  in  vacuo,  deposits  reddish  prismatic 
deliquescent  crystals  having  the  composition  K2S,5H2O. 

When  potassium  carbonate  and  sulphur  are  heated  together  a 
mixture  of  the  higher  sulphides  of  potassium  with  potassium  thio- 
sulphate  is  obtained,  thus  — 


3K2CO3+   8S  =  2K2S3  +  K2S2 

3K2CO3  +  12S  =  2K2S5  +  K2S2O3  +  3CO2. 

The    reddish  -  brown   solid   product   was   named  by  the   early 
chemists  hepar  sulphuris,  or  "  liver  of  sulphur." 


SODIUM. 

Symbol,  Na  =  23.00. 

Occurrence. — The  most  abundant  natural  compound  of  sodium 
is  the  chloride,  which  is  present  in  sea-water  and  in  many  salt 
lakes  and  springs.  Enormous  deposits  of  sodium  chloride  or 
rock-salt  are  found  in  Cheshire,  Lancashire,  and  other  parts  of 
the  world.  As  nitrate,  this  element  occurs  in  large  quantities  in 
Chili  and  Peru,  and  in  combination  with  silicic  acid  it  is  a  con- 
stituent of  many  rocks. 

Modes  Of  Formation.— Sodium  was  first  isolated  by  Davy,  by 
the  electrolysis  of  sodium  hydroxide.  On  a  manufacturing  scale 
it  has  been  obtained  by  the  various  processes  described  under 
potassium,  the  history  of  the  commercial  preparation  of  these  two 
elements  being  practically  identical.  Sodium,  however,  does  not 
form  any  explosive  compound  with  carbon  monoxide,  so  that  the 
manufacture  in  this  case  has  been  free  from  this  difficulty. 

At  the  present  time  sodium  (and  also  potassium)  is  almost  ex- 
clusively  obtained  by  electrolytic  methods. 

(i.)  Castner's  Process. — This  method  consists  in  the  electrolysis 
of  fused  sodium  hydroxide,  and  is,  in  fact,  simply  the  original 
process  by  which  Davy  first  obtained  the  metal  adapted  to  modern 
resources  of  electrical  power.  The  apparatus  employed  is  shown 
in  section  in  Fig.  134.  The  caustic  soda  is  contained  in  an  iron 


Sodium 


527 


pot  P,  set  in  suitable  brick-work,  and  is  kept  in  a  melted  state  by 
a  ring  of  gas  flames  below.  Through  the  bottom  of  this  vessel 
passes  the  cathode,  which  is  maintained  steady  in  its  position  by 
the  caustic  soda  in  the  lower  and  narrow  part  of  vessel  P  being  in 
the  solidified  state.  The  anodes  A  are  suspended  from  above 
round  the  cathode,  and  are  prevented  from  touching  it  by  means 
of  a  wirework  cylinder  which  hangs  from  the  vessel  V.  This  vessel 
is  an  iron  cylinder  having  a  lid  at  the  top,  and  is  the  receiver  in 
which  the  sodium  collects. 

The  products  of  the  electrolysis  are  oxygen,  hydrogen,  and 
sodium.  The  oxygen  liberated  at  the  anodes  escapes  by  the 
opening  O  in  the  lid.  The 
sodium  floats  up  to  the  sur- 
face of  the  molten  caustic 
in  the  receiver  V,  and  is 
withdrawn  from  time  to 
time  by  means  of  a  per- 
forated ladle,  which  allows 
the  caustic  to  drain  through, 
but  holds  the  liquid  metal, 
owing  to  the  extremely  high 
surface  tension  of  the  latter. 
The  hydrogen  which  is  also 
liberated  at  the  cathode 
escapes  through  the  loosely 
fitting  lid  of  the  receiver.* 

(2.)  Borcherf  Process. — 
It  will  be  evident  from  an 

economic  (and  therefore  the  manufacturer's)  point  of  view  that  the 
hydrogen  liberated  in  the  above  process  represents  wasted  electrical 
energy.  Many  attempts,  therefore,  have  been  made  to  substitute 
fused  sodium  chloride  for  the  hydroxide.  The  practical  difficulties  to 
be  overcome  in  this  case  are  more  serious,  owing  partly  to  the  higher 
temperature  required,  and  also  to  the  corrosive  action  exerted  by 
the  fused  chloride  upon  the  materials  of  which  the  vessels  are 
constructed.  On  the  other  hand,  it  will  be  evident  that  both  the 
products  of  the  electrolysis  in  this  case  will  have  commercial 
value.  Borchers'  apparatus  is  shown  in  section  in  Fig.  135.  It 


FIG.  134. 


*  This  process  is  extensively  employed  at  Oldbury,  near  Birmingham,  and 
at  the  works  of  the  Niagara  Electrical  Company. 


528 


Inorganic  Chemistry 


consists  essentially  of  a  U-tube  made  in  two  parts,  the  wide  limb 
being  of  fireclay  and  the  narrow  part  of  iron,  the  two  parts  being 
clamped  together.  To  prevent  leakage  at  the  joint,  a  square 
tube  W  is  interposed  between  the  two  parts,  which  is  kept  cool  by 
a  flow  of  water  through  it.  This  causes  the  solidification  of  the 
sodium  chloride  in  the  form  of  a  layer  all  round  the  junction. 
Chlorine  is  liberated  at  the  anode  and  escapes  by  the  pipe  P. 
The  narrow  limb  C  is  itself  the  cathode,  and  the  sodium  there 


FIG.  135. 

produced  overflows  down  the  side  pipe  into  a  suitable  receiver. 
Fresh  sodium  chloride  is  added  as  required  through  the  tube  D. 

Properties. — Sodium  closely  resembles  potassium  in  its  general 
properties.  It  is  a  soft,  white  metal  which  can  be  readily  moulded 
by  the  fingers,  and  is  easily  pressed  into  wire.  At  -  20°  it  is  hard. 
The  colour  of  sodium  vapour  is  violet,  while  the  colour  exhibited 
by  a  thin  film  of  the  metal,  obtained  by  sublimation  in  vacuo,  is 
greenish-blue.  The  vapour-density  of  sodium  is  about  12  (Dewar 
and  Scott),  showing  that  this  metal  in  the  vaporous  state  is 
monatomic. 

Like  potassium,  sodium  dissolves  in  liquid  ammonia,  yielding  a 


Sodium  Peroxide  529 

blue  solution.  When  heated  in  the  air,  sodium  burns,  forming 
the  peroxide,  Na2O2.  Perfectly  dry  air  or  oxygen  is  without 
action  upon  the  metal. 

When  heated  in  hydrogen,  sodium  forms  the  hydride,  NaH^ 
analogous  to  the  potassium  compound,  but  not  spontaneously 
inflammable  in  air.  When  this  is  heated  to  about  300°  in  vacuo 
the  whole  of  the  hydrogen  is  evolved. 

Alloy  Of  Sodium  and  Potassium.— When  these  two  metals  are 
melted  together  beneath  petroleum  an  alloy  is  obtained  which  is 
liquid  at  ordinary  temperatures.  When  prepared  and  preserved 
out  of  contact  with  air  the  alloy  resembles  mercury  in  appearance. 
This  alloy  is  employed  in  the  construction  of  thermometers  for 
registering  high  temperatures,  where  mercury  would  be  inad- 
missible. 

Oxides  Of  Sodium.  —  Two  oxides  are  known,  viz.,  sodium 
monoxide,  Na2O,  and  sodium  dioxide,  or  peroxide,  Na2O2.  Of 
these  the  peroxide  is  the  more  import  ant. 

Sodium  Oxide  (or  sodium  monoxide),  Na2O,  is  produced  as  a 
white  amorphous  compound  when  sodium  is  partially  oxidised  in 
a  limited  supply  of  oxygen,  and  the  excess  of  the  metal  subse- 
quently removed  by  distillation  in  vacuo.* 

Sodium  Peroxide,  Na2O2,  is  obtained  by  allowing  sodium  to 
burn  briskly  in  oxygen.  It  is  a  yellowish-white  solid,  which  de- 
composes in  contact  with  water,  with  considerable  rise  of  tem- 
perature and  evolution  of  oxygen — 

Na202+H2O  =  2NaHO  +  O. 

The  oxygen  which  is  evolved  contains  appreciable  quantities  of 
ozone.  When  sodium  peroxide  is  slowly  added  to  water  or  to  dilute 
hydrochloric  acid  in  the  cold,  hydrogen  peroxide  is  formed — 

Na2O2  +  2H2O  =  2NaH  O  +  H2O2. 

Owing  to  the  readiness  with  which  it  gives  up  oxygen,  sodium 
peroxide  is  a  powerful  oxidising  agent,  and  as  such  finds  many 
uses  in  the  laboratory.  Thus  it  readily  converts  chromic  com- 
pounds into  chromates. 

Sodium  peroxide  forms  a  crystalline  hydrate  of  the  composition, 
Na2O2,8H2O  (page  228).  When  heated  in  either  nitrous  or  nitric 

*  Rengade,  Compt.  Rend.,  1906. 

2  L 


530 


Inorganic  Chemistry 


oxides  it  yields  sodium  nitrite  ;  in  the  former  case  with  the  elimina- 
tion of  nitrogen  — 


Na2O2-f2N2O  = 

Na2O2  +  2NO  =2NaNO2. 

Sodium  Hydroxide  (caustic  soda),  NaHO.—  This  compound  is 
produced  when  sodium  is  brought  into  contact  with  water,  and  also 
when  either  sodium  monoxide  or  peroxide  is  dissolved  in  water. 
On  the  large  scale  caustic  soda  is  prepared  by  the  action  of  lime 
upon  a  boiling  solution  of  sodium  carbonate  (see  Caustic  Potash). 

The  so-called  tank  liquors  (obtained  in  the  manufacture  of  sodium 
carbonate  by  the  Leblanc  process,  q.v.)  are  heated  to  the  boiling- 
point,  and  an  excess  of  lime  is  stirred  into  the  mixture.  The 
sodium  sulphide  present  in  the  tank  liquor  is  oxidised  into  sulphate 


FIG.  136. 

•by  the  combined  action  of  air  injected  into  the  mixture,  and  of 
sodium  nitrate,  which  is  added  for  this  purpose.  The  liquor,  after 
being  causticised,  is  decanted  or  filtered  from  the  precipitated 
calcium  carbonate,  and  is  concentrated  in  large  cast-iron  hemi- 
spherical pans.  The  decomposition  suffered  by  the  sodium  nitrate 
depends  upon  the  temperature  and  concentration  of  the  liquid  ;  at 
300°  to  360°  the  change  may  be  expressed  by  the  equation  — 


The  liberated  oxygen  oxidises  the  sulphides  to  sulphates. 

Caustic  soda  is  now  being  manufactured  by  the  electrolysis  of 
brine.  The  apparatus  devised  by  Castner  for  this  purpose  is  seen 
in  Fig.  136.  It  consists  of  a  rectangular  vessel  divided  into  three 
compartments.  Upon  the  floor  of  the  vessel  there  is  a  layer  of 


Sodium  Chloride  531 

mercury  about  £th  of  an  inch  deep.  The  partitions,  which  are 
non-porous,  dip  into  narrow  gutters  across  the  bottom,  but  do  not 
actually  touch  the  bottom,  so  that  when  the  tank  is  gently  oscil- 
lated the  mercury  can  flow  from  one  compartment  to  the  other, 
while  the  liquid  above  is  prevented  from  so  doing.  The  two 
outside  compartments  are  filled  with  brine,  while  the  centre  one 
contains  water  ;  and  in  this  is  placed  the  cathode,  consisting  of  a 
number  of  metal  plates.  Since  the  partitions  are  non-porous  the 
current  will  pass  from  the  carbon  anodes  through  the  salt  solution 
to  the  mercury,  which  in  the  two  extreme  compartments  then 
becomes  the  cathode.  It  then  passes  from  the  mercury  in  the 
middle  space,  which  now  becomes  the  anode  of  this  cell,  through 
the  aqueous  liquid  to  the  metal  cathode  which  is  there  suspended. 
In  the  outside  compartments  the  sodium  chloride  is  electrolysed  ; 
the  chlorine  discharged  at  the  carbon  anodes  escapes  by  the  pipes 
P  P,  while  the  sodium  dissolves  in  the  mercury  cathodes.  During 
the  process  a  slow  rocking  movement  is  given  to  the  tank  by  means 
of  the  excentric  represented  at  E,  whereby  the  mercury  is  caused 
to  flow  to  and  fro  along  the  bottom.  In  the  middle  compartment 
the  sodium  contained  in  the  amalgam  is  transported  to  the  cathode, 
where  it  dissolves  in  the  water,  forming  sodium  hydroxide. 

Sodium  hydroxide  is  a  white,  strongly  caustic,  and  highly  de- 
liquescent solid.  It  is  soluble  in  water,  with  considerable  rise  of 
temperature,  and  a  concentrated  aqueous  solution  when  cooled 
to  —8°,  deposits  a  crystalline  hydrate,  having  the  composition 
2NaHO,7H2O. 

Sodium  Chloride,  NaCl.— Of  the  compounds  of  sodium  with 
the  halogens  the  chloride  is  the  most  important.  In  warm 
climates,  as  upon  the  shores  of  the  Mediterranean,  sodium  chloride 
is  obtained  by  the  evaporation  of  sea-water  in  large  shallow  basins 
or  pools,  constructed  upon  the  sea-shore  and  exposed  to  the  sun's 
heat.  As  the  brine  concentrates  in  these  salterns,  the  crystals 
of  salt  are  raked  off  the  liquid  and  allowed  to  drain  in  heaps  at 
the  side  of  the  pools.  The  mother-liquors,  known  as  bittern^ 
were  formerly  utilised  for  the  extraction  of  the  bromine  which  they 
contam. 

Salt  is  obtained  from  salt-beds,  where  it  is  found  in  enormous 
deposits,  either  by  direct  mining  operations,  when  the  salt  is 
sufficiently  pure,  or  by  first  dissolving  the  material  in  water, 
whereby  insoluble  admixed  impurities  are  removed,  and  afterwards 
evaporating  the  brine  so  obtained.  The  latter  method  is  carried 


532 


Inorganic  Chemistry 


out  by  sinking  borings  through  the  upper  strata  of  rock,  and 
sending  water  down  to  the  salt-beds  beneath.  The  brine  is  then 
pumped  up  and  the  salt  obtained  by  evaporation.  The  first  stage 
of  the  concentrating  process,  especially  where  the  brine  is  not  very 
strong,  is  in  some  parts  .carried  on  by  exposing  the  liquid  to  the 


FIG.  137. 

wind.  This  is  effected  by  causing  the  solution  to  trickle  over 
erections  of  brushwood  known  as  graduators  (Fig.  137),  which  are 
built  so  that  the  prevailing  winds  blow  across  them.  The  brine  is 
pumped  up  into  the  wooden  troughs  running  along  the  top,  from 
which  it  escapes  by  a  number  of  openings,  a,  a,  a,  and  flows  over 
the  pile  of  brushwood  down  into  the  reservoir  upon  which  the 


Sodium  Chloride 


533 


erection  is  constructed.  In  this  way  the  solution  is  made  to 
expose  a  large  surface  to  the  air,  and  it  quickly  reaches  a  concen- 
tration when  it  contains  20 
per  cent,  of  salt  in  the 
solution.  The  liquor  is  then 
evaporated  in  shallow  iron 
pans  by  means  of  artificial 
heat,  and  as  the  salt  crys- 
tallises it  is  lifted  out  by 
means  of  perforated  iron 
skimmers.  Salt  obtained 
in  this  manner  always  con- 
tains small  quantities  of 
other  salts,  such  as  sodium 
sulphate,  calcium  sulphate, 
calcium  and  magnesium 
chlorides.  The  presence 
of  chlorides  of  magnesium 
or  calcium  causes  the  salt 
to  become  moist,  especially 
in  damp  weather. 

Pure  sodium  chloride 
may  be  prepared  by  add- 
ing hydrochloric  acid  to  a 
strong  aqueous  solution  of 
salt ;  the  sodium  chloride  is 
thereby  precipitated,  while 
the  other  salts  remain  in 
solution. 

Sodium  chloride  forms 
colourless,  cubical  crystals, 
which  are  anhydrous.  If 
deposited  at  -  10°  it  crystal- 
lises in  monosymmetric 
prisms,  with  two  molecules 
of  water  of  crystallisation, 
which  at  the  ordinary  temperature  lose  their  water  and  break  up 
into  minute  cubes. 

Sodium  chloride  is  a  necessary  article  of  food  for  man  and  other 
animals  ;  it  is  estimated  that  about  20  Ibs.  of  salt  per  head  of 
population  is  annually  used,  directly  or  indirectly,  for  this  purpose. 


534  Inorganic  Chemistry 

The  hydrochloric  acid  present  in  the  gastric  and  other  acid  fluids 
of  the  stomach  is  derived  from  the  decomposition  of  sodium  chloride 
which  is  taken  into  the  organism. 

Enormous  quantities  of  sodium  chloride  are  employed  in  the 
alkali  industry,  and  all  the  chlorine  that  is  manufactured  is  derived 
primarily  from  this  compound. 

Sodium  Bromide,  NaBr,  and  Sodium  Iodide,  Nal,  are  pre- 
pared by  methods  similar  to  those  for  obtaining  the  potassium 
compounds.  They  are  both  isomorphous  with  sodium  chloride, 
and  when  deposited  at  low  temperatures  they  form  monosymmetric 
crystals  containing  two  molecules  of  water. 

Sodium  Carbonate,  Na2CO3.—  The  preparation  of  this  com- 
pound is  carried  on  by  three  methods,  and  constitutes  that  important 
industry,  the  alkali  manufacture.  Two  of  these  processes  are  known 
by  the  names  of  their  respective  discoverers,  namely,  the  Leblanc 
process  and  the  Solvay  process,  the  latter  being  also  known  as 
the  ammonia-soda  process.  The  third  is  a  modern  electrolytic 
method. 

I.  The  Leblanc  method  of  manufacture  consists  essentially  of 
three  processes,  namely  — 

(i.)  The  conversion  of  sodium  chloride  into  sodium  sulphate 
by  the  action  of  sulphuric  acid,  known  as  the  salt-cake 
process.  Two  chemical  reactions  are  involved  in  the 
process  — 

NaCl  +  H2SO4      =NaHSO4  +  HCl. 
NaCl  +  NaHSO4=Na2SO4   +HC1. 

(2.)  The  decomposition  of  sodium  sulphate,  salt-cake^  by  means 
of  calcium  carbonate  (limestone)  and  coal,  at  a  high 
temperature,  whereby  a  crude  mixture  of  sodium  car- 
bonate and  calcium  sulphide  is  obtained,  known  as 
black-ash.  This  black-ash  process  takes  place  in  accord- 
ance with  the  following  equation  — 


The  change  may  be  conveniently  regarded  as  taking  place  in  two 
stages,  which  proceed  simultaneously  according  to  the  equations  — 

Na2SO4+2C  =  Na2S  +  2CO2. 
Na2S  +  CaCO3  =  CaS  +  Na2CO3. 

(3.)  The  process  of  extracting  and  purifying  the  sodium  car- 
bonate contained  in  the  black-ash. 


Sodium  Carbonate 


535 


(i.)  The  Salt-cake  Process.— The  first  stage  of  this  process  is 
usually  carried  on  in  a  large  cast- 
iron  pan  (D,  Fig.  138),  built  into 
a  furnace  in  such  a  manner  that 
it  shall  be  heated  as  uniformly  as 
possible.  The  charge  of  common 
salt  is  placed  in  the  covered  pan, 
and  the  requisite  quantity  of  sul- 
phuric acid  is  then  run  in.  Hydro- 
chloric acid  is  given  off  in  tor- 
rents, according  to  the  first  of  the 
above  equations,  and  the  gas  is 
led  away  by  the  pipe  E  in  the 
arched  roof  to  the  condensing- 
towers,  where  it  is  absorbed  by 
water  (see  Hydrochloric  Acid, 
page  369).  The  mixture  is  heated 
until  it  begins  to  stiffen  into  a 
solid  mass,  when  the  damper  h 
is  raised  and  the  mass  is  raked 
out  of  the  pan  on  to  the  hearth 
of  the  roaster  or  reverberatory 
furnace,  b.  Here  it  is  exposed 
to  the  hot  gases  from  the  coke  fire 
a,  which  sweep  over  it  and  ulti- 
mately raise  its  temperature  nearly 
to  a  red  heat,  whereby  the  second 
of  the  above  reactions  is  com- 
pleted. The  acid  gas,  together 
with  the  fire  gases,  leave  the 
roaster  by  the  chimney  e,  and  are 
also  led  to  condensing  -  towers, 
where  the  hydrochloric  acid  is 
absorbed.  The  mass  is  from  time 
to  time  raked  or  worked  by  means 
of  side -openings  or  "working 
doors  "  in  the  roaster,  and  as  soon 
as  the  operation  is  completed  the 
salt-cake  is  withdrawn.  The  salt- 
cake  so  obtained  usually  contains 
from  95  to  q.6  per  cent,  of  normal 


536 


Inorganic  Chemistry 


sodium  sulphate,  Na2SO4  ;  the  remaining  4  or  5  per  cent,  consist- 
ing of  hydrogen  sodium  sulphate,  NaHSO4,  undecomposed  sodium 

chloride,  and  such  impurities  as  were 
originally  present  in  the  salt. 

(2.}  The  Black-ash  Process.— The 
salt-cake  is  mixed  with  limestone  (or 
chalk)  and  coal  dust  (slack\  and 
heated  in  a  reverberatory  furnace 
known  as  the  black-ash  or  balling' 
furnace.  As  the  mixture  softens  with 
the  heat  it  requires  to  be  thoroughly 
mixed  together,  which,  in  the  older 
forms  of  furnace  (still  used  in  many 
places),  is  accomplished  by  manual 
labour.  Fig.  139  shows  such  a  furnace 
in  section.  The  materials  are  intro- 
duced by  the  hopper  k  on  to  the 
hearth  /,  where  they  are  exposed  to 
the  hot  gases  from  the  fire  a  ;  and 
as  the  decomposition  proceeds  they 
are  raked  along  to  the  more  strongly- 
heated  front  portion  of  the  hearth  h. 
During  this  process  carbon  dioxide 
is  freely  evolved,  the  escaping  bubbles 
of  gas  giving  the  semi-fluid  mass 
the  appearance  of  boiling.  As  the 
temperature  rises  and  the  process 
approaches  completion,  the  mass 
thickens,  when  it  is  worked  up  into 
large  balls  by  means  of  rakes  or 
paddles.  At  this  stage  carbon  mon- 
oxide begins  to  be  evolved,  the  bubbles 
of  which,  bursting  from  the  doughy 
material,  become  ignited  and  burn 
upon  its  surface  as  small  jets  of  flame 
coloured  yellow  by  the  soda.  As  soon 
as  these  appear  the  ball  is  quickly  withdrawn  from  the  furnace. 
The  formation  of  carbon  monoxide  at  the  high  temperature  reached 
at  this  point  in  the  process  is  due  to  the  action  of  carbon  upon  the 
limestone  according  to  the  equation — 

CaCO3  +  C  =  CaO  +  2CO, 


Sodium  Carbonate  537 

excess  of  these  materials  being  intentionally  present  in  the  mixture. 
The  effect  of  the  escaping  carbon  monoxide  at  this  point  in  the 
process,  in  rendering  the  black-ash  light  and  porous  (an  important 
consideration  in  view  of  the  next  operation),  is  similar  to  that  of 
baking-powder  when  used  for  cooking  purposes.  The  heated 
gases  from  the  furnace  are  made  to  pass  over  large  evaporating 
pans,  /*,  where  liquors  from  a  subsequent  process  are  concen- 
trated. 

In  the  more  modern  forms  of  black-ash  furnace,  the  mixing  and 
working  up  of  the  materials  is  accomplished  mechanically  by 
means  of  a  revolving  hearth.  Fig.  140  shows  the  general  arrange- 
ment of  a  revolving  black-ash  furnace.  The  mixture  is  placed  in 
the  cylinder  *,  which  is  made  to  slowly  revolve  upon  its  horizontal 
axis.  The  heated  gases  from  the  fire  a  pass  through  this  revolv- 
ing hearth  ;  they  are  then  conveyed  through  a  dust-chamber,  m,  and 
finally  over  concentrating-pans.  Limestone  and  two-thirds  of  the 
coal  are  first  thrown  into  the  furnace  and  heated  until  the  blue 
flame  of  burning  carbon  monoxide  makes  its  appearance,  when 
the  salt-cake  along  with  the  rest  of  the  coal  is  added,  and  the 
process  continued  until  the  yellow  flames  appear  upon  the  surface 
of  the  mass.  The  contents  of  the  cylinder  are  then  thrown  out 
into  iron  trucks  beneath. 

Black-ash  is  a  mixture  of  variable  composition,  containing — 

Sodium  carbonate,  Na2CO3    .  from  40  to  45  per  cent. 

Calcium  sulphide,  CaS  .  „       30  „  33         „ 

Calcium  carbonate,  CaCO3     .  „         6  „  10         „ 

Coke „         4  „     7         „ 

Calcium  oxide,  CaO        .  „         2  „     6        „ 

And  smaller  quantities  of  sodium  chloride,  sodium  sulphate,  sodium 
sulphite,  sodium  sulphide,  sodium  thiosulphate,  oxides  of  iron 
and  alumina. 

(3.)  Lixiviation  of  Black-ask, — The  lixiviation  of  black-ash  is 
carried  on  in  a  series  of  tanks,  so  arranged  that  the  liquid  can  be 
made  to  pass  from  one  to  the  other.  The  action  of  water  upon 
the  black-ash  is  more  than  a  simple  process  of  dissolving  the 
sodium  carbonate  from  the  mixture,  for  in  the  presence  of  water 
chemical  action  takes  place  between  some  of  the  ingredients. 
Thus  the  lime  reacts  upon  sodium  carbonate,  forming  sodium 
hydroxide,  hence  the  tank  liqtior  always  contains  caustic  soda  in 
varying  quantities.  Under  certain  conditions  of  temperature  and 


538  Inorganic  Chemistry 

dilution,  the  calcium  sulphide  also  reacts  upon  the  sodium  car- 
bonate, forming  sodium  sulphide  and  calcium  carbonate,  thus  — 

CaS  +  Na2CO3  =  CaCO3  +  Na2S. 

Also  by  the  oxidising  influence  of  atmospheric  oxygen,  calcium 
sulphide,  CaS,  is  converted  into  calcium  sulphate,  CaSO4,  which 
in  its  turn  is  acted  upon  by  the  sodium  carbonate,  involving  loss 
of  this  product  — 


The  process  of  lixiviation  is  carried  on  as  rapidly  as  possible, 
and  at  temperatures  ranging  from  about  30°  (for  the  dilute  liquors) 
to  about  60°  (for  those  more  concentrated)  ;  for  the  formation  of 
sodium  sulphide  diminishes  as  the  concentration  of  the  liquid 
increases.  The  tank  liquor,  after  settling,  is  then  either  at  once 
concentrated  by  evaporation,  when  the  soda  crystallises  out,  leav- 
ing the  caustic  soda  in  the  mother-liquor,  or  it  is  submitted  to  the 
action  of  carbon  dioxide,  whereby  both  the  caustic  soda  and  the 
sodium  sulphide  are  converted  into  sodium  carbonate,  thus  — 

2NaHO  -f  CO2=  Na2CO3  +  H2O. 
Na2S  +  CO2  +  H2O  =  Na2CO3  +  H2S. 

The  concentration  of  the  tank  liquor  is  accomplished  in  the 
shallow  pans  above  mentioned,  by  means  of  the  waste  heat  from 
the  black-ash  furnace  ;  and  the  product  obtained  by  evaporating  the 
liquid  is  usually  calcined  at  a  red  heat  in  an  ordinary  reverberatory 
furnace.  This  substance  is  known  as  soda-ash^  and  when  dissolved 
in  water,  and  the  solution  allowed  to  crystallise,  the  so-called  soda 
.crystals  are  obtained,  having  the  composition  Na2CO3,10H2O. 

II.  The  Ammonia-  Soda  Process.  —  This  process  is  based  upon  the 
fact,  that  hydrogen  ammonium  carbonate  (bicarbonate  of  ammonia) 
is  decomposed  by  a  strong  solution  of  sodium  chloride,  according 
to  the  equation  — 

H(NH4)CO3  +  NaCl  =  HNaCO3  +  NH4Cl. 

In  practice  the  brine  is  first  saturated  with  ammonia  gas,  and 
the  cooled  ammoniacal  liquid  is  then  charged  with  carbon  dioxide, 
under  moderate  pressure,  in  carbonating  towers. 

The  hydrogen  sodium  carbonate  (bicarbonate  of  soda),  being 
much  less  soluble,  separates  out,  leaving  the  more  soluble  am- 


Sodium  Carbonate  539 

monium  chloride  in  solution,  from  which  the  ammonia  is  recovered 
by  subsequent  treatment  with  lime. 

The  hydrogen  sodium  carbonate  is  converted  into  normal  sodium 
carbonate  by  calcination,  and  the  carbon  dioxide  evolved  is  again 
utilised  in  carbonating  a  further  quantity  of  ammoniacal  brine  — 


These  two  processes,  namely,  the  Leblanc  and  the  ammonia-soda 
process,  have  been  keen  competitors  for  a  number  of  years  ;  and 
a  glance  at  the  figures  giving  the  annual  output  from  the  two 
sources  shows  how  rapidly  and  steadily  the  younger  process  has 
gained  upon  its  older  rival.  Indeed,  there  can  be  little  doubt  that 
but  for  the  value  of  the  hydrochloric  acid  which  is  simultaneously 
produced  in  the  Leblanc  process,  this  method  would  before  now 
have  ceased  to  exist  as  a  manufacture.  Now,  however,  both  of 
these  processes  are  threatened  by  the  advent  of  a  new  and  formid- 
able rival  in  the  electrolytic  method. 

III.  The  Electrolytic  Process  (Hargreaves-Bird).  —  In  this 
method  a  solution  of  sodium  chloride  (brine,  pumped  direct  from 
the  salt-beds)  is  submitted  to  electrolysis  in  a  cell  of  special  con- 
struction. This  consists  of  an  oblong  box  divided  longitudinally 
into  three  compartments,  the  centre  one  being  comparatively 
large,  while  the  two  extreme  compartments  are  quite  narrow.  The 
partitions  which  divide  the  box  in  this  manner  are  made  of  a 
"  composition  "  consisting  largely  of  asbestos  ;  and  are  of  such  a 
nature  that  when  the  middle  compartment  is  filled  with  brine. 
none  of  the  liquid  percolates  or  oozes  through  into  the  side  cham- 
bers. These  asbestos  diaphragms  are  backed  on  their  outer  sides 
by  a  network  of  copper  wire  which  is  made  the  cathode  in  the 
system.  The  anode  consists  of  pieces  of  gas-carbon  which  are 
suspended  in  the  brine  in  the  centre  chamber.  Although  the 
asbestos  diaphragms  are  water-tight,  in  the  sense  that  they  do  not 
allow  the  brine  to  pass  from  the  middle  to  the  outer  compartments, 
they  are  nevertheless  sufficiently  porous  to  keep  the  copper  wire 
cathodes  moist,  and  to  allow  therefore  of  the  passage  of  the 
current.  Chlorine  is  evolved  at  the  anode,  and  is  conveyed  away 
directly  to  lime  chambers  and  converted  into  bleach  ing-powder. 
The  sodium  ions  pass  freely  through  the  asbestos  partitions  to  the 
cathodes,  there  generating  sodium  hydroxide  ;  while  a  stream  of 
steam  and  carbon  dioxide  which  is  passed  through  the  narrow 


540  Inorganic  Chemistry 

compartments  immediately  converts  the  hydroxide  into  carbonate, 
which  is  thus  washed  away  from  the  cathodes  as  fast  as  it  is 
formed.  The  solution  so  obtained  is  sufficiently  concentrated  to 
deposit  crystals  of  sodium  carbonate  on  cooling. 

The  "  soda"  obtained  by  this  process,  which  is  now  being  carried 
out  on  an  extensive  scale  at  Middlewich,  Cheshire,  is  extremely 
pure,  containing  from  97  to  98  per  cent,  of  sodium  carbonate,  and 
only  about  I  per  cent,  of  sodium  chloride. 

Sodium  carbonate  crystallises  in  large,  transparent,  monosym- 
metric  crystals,  commonly  known  as  "  soda  "  or  "  washing-soda," 
having  the  composition  Na2CO3,10H2O.  On  exposure  to  the  air 
the  crystals  give  up  water,  and  become  effloresced  upon  the  surface, 
and  finally  fall  to  powder,  having  the  composition  Na2CO3,H2O. 
When  crystallised  from  hot  solutions,  it  forms  rhombic  crystals, 
containing  7H2O.  The  solubility  of  sodium  carbonate  in  water 
increases  with  rise  of  temperature,  reaching  a  maximum  at  32.5°, 
when  loo  parts  of  water  dissolve  59  parts  of  the  salt.  Above  this 
temperature  the  solubility  falls,  and  at  100°  the  amount  dissolved 
is  45.4  parts. 

Hydrogen  Sodium  Carbonate  (bicarbonate  of  soda),  HNaCO3, 
may  be  obtained  by  the  action  of  carbon  dioxide  upon  the  normal 
carbonate,  either  in  solution  or  as  crystals — 

Na2CO3,10H2O  +  CO2  =  2HNaCO3  +  9H2O. 

The  greater  part  of  the  bicarbonate  of  soda  of  commerce  is 
obtained  in  the  ammonia-soda  process  above  described. 

This  salt  is  less  soluble  in  water  than  the  normal  carbonate. 
Thus,  100  parts  of  water  at  different  temperatures  dissolve  the 
following  quantities  of  these  compounds — 

10°.  20°.  30°.  40°. 

Na2CO3.     .     .     12.6  21.4  38.1  50    parts. 

HNaCO3    .     .       8.8  9.8  10.8  11.7     „ 

When  a  solution  of  hydrogen  sodium  carbonate  is  heated,  the  salt 
gives  off  a  portion  of  its  carbon  dioxide,  and  on  cooling  the  solution 
deposits  crystals  having  the  composition  Na2CO3,2HNaCO3,2H2O, 
known  as  sodium  sesquicarbonate.  On  continued  boiling,  the  salt 
is  completely  converted  into  the  normal  carbonate.  Sodium 
sesquicarbonate  occurs  as  a  natural  deposit  in  Egypt,  Africa, 
South  America,  and  elsewhere,  known  as  trona>  from  which  the 
name  natrium  is  derived. 


Sodium  Nitrate  54  r 

Sodium  Sulphate  (Glauber's  salt\  Na2SO4,10H2O,  occurs  native 
in  the  anhydrous  condition  as  the  mineral  thenardite,  and  as  a 
double  sulphate  of  sodium  and  calcium,  Na2SO4,CaSO4,  in  the 
mineral  Glauberite. 

It  is  manufactured  in  immense  quantities  in  the  first  (salt-cake) 
process  in  the  alkali  manufacture,  by  the  Leblanc  method. 

It  is  also  obtained  in  large  supplies  from  the  Stassfurt  deposits, 
by  double  decomposition  between  magnesium  sulphate  (from 
kieserite]  and  sodium  chloride. 

The  solution  of  the  mixed  salts,  when  cooled  a  few  degrees 
below  o°,  deposits  sodium  sulphate,  and  the  soluble  magnesium 
chloride  remains  in  solution  — 

2NaCl  +  MgSO4  =  Na2SO4+MgCl2. 

Sodium  sulphate  is  also  manufactured  by  the  action  of  sulphur 
dioxide  and  oxygen  upon  sodium  chloride.  This  is  known  as 
HargreavJs  process.  The  reaction  is  expressed  by  the  equation  — 


This  process  is,  in  essence,  the  production  of  sodium  sulphate 
from  sodium  chloride  and  the  constituents  of  sulphuric  acid,  with- 
out the  intermediate  manufacture  of  the  acid.  The  gases  from 
pyrites  burners,  similar  to  those  used  by  the  "vitriol"  manufacturer, 
together  with  steam,  are  passed  through  a  series  of  cast-iron 
cylinders  containing  sodium  chloride,  and  maintained  at  a  tem- 
perature of  500°  to  550°.  Many  days  are  required  for  the  com- 
plete conversion  of  the  chloride  into  sulphate  by  this  process. 

Sodium  sulphate  crystallises  in  colourless  prisms  belonging  to 
the  monosymmetric  system,  containing  ten'  molecules  of  water  ; 
when  exposed  to  the  air  the  crystals  effloresce,  and  when'  heated 
to  33°  they  melt  in  their  own  water  of  crystallisation  (see  page 

153). 

When  sodium  sulphate  is  heated  with  sulphuric  acid,  in  the  pro- 
portions required  by  the  following  equation,  hydrogen  sodium 
sulphate  is  formed  — 

Na2S04  +  H2S04  =  2HNaS04. 

Sodium  Nitrate,  NaNO3,  occurs  associated  with  other  salts  in 
Bolivia  and  Peru,  as  cubical  nitre,  or  Chili  saltpetre.  The  crude 


542  Inorganic  Chemistry 

salt  is  purified  by  solution  in  water,  and  crystallisation.  It  forms 
rhombohedral  crystals,  isomorphous  with  calcspar. 

Sodium  nitrate  is  very  soluble  in  water.  100  parts  of  water  dis- 
solve at  o°,  68.8  parts  ;  at  40°,  102  parts  ;  and  at  100°,  180  parts  of 
the  salt.  When  exposed  to  the  air,  the  salt  absorbs  moisture,  and 
on  this  account  cannot  be  employed  as  a  substitute  for  potassium 
nitrate  in  the  manufacture  of  gunpowder,  or  in  pyrotechny.  Its 
chief  uses  are  for  the  manufacture  of  nitric  acid  ;  for  the  manufacture 
of  potassium  nitrate  by  double  decomposition  with  potassium 
chloride  ;  and  as  an  ingredient  in  artificial  manures. 

Sodium  Phosphates.  —  The  most  important  of  these  compounds 
is  the  hydrogen  disodium  orthophosphate,  or  common  phosphate 
of  soda,)  HNa2PO4.  This  salt  is  prepared  on  a  large  scale,  by 
adding  sodium  carbonate  to  phosphoric  acid  until  the  solution  is 
alkaline,  and  then  filtering  and  evaporating  the  solution,  when 
large  transparent  prisms,  belonging  to  the  monosymmetric  system, 
are  deposited,  having  the  composition  HNa2PO4,12H2O.  Exposed 
to  the  air  the  crystals  effloresce,  and  when  heated  become  an- 
hydrous. The  salt  melts  at  35°. 

100  parts  of  water  at  10°  dissolve  4.1  parts  ;  at  50°,  43.3  parts  ; 
and  at  100°,  108.2  parts  of  the  anhydrous  salt. 

Normal  Sodium  Orthophosphate,  Na3PO4,  is  obtained  from 
hydrogen  disodium  phosphate,  by  evaporating  a  solution  of  the 
latter  salt  with  sodium  hydroxide,  until  the  liquid  crystallises  — 

HNa2PO4  +  NaHO  =  Na3PO4+H2O. 

This  salt  contains  twelve  molecules  of  water,  and  forms  thin 
six-sided  prisms.  Its  aqueous  solution  is  strongly  alkaline,  and 
absorbs  atmospheric  carbon  dioxide,  with  the  formation  of  hydrogen 
sodium.  carbonate  and  hydrogen  disodium  phosphate,  thus  — 


Dihydrogen  Sodium  Orthophosphate,  H2NaPO4,  is  obtained 
when  phosphoric  acid  is  added  to  ordinary  phosphate  of  soda,  until 
the  liquid  gives  no  precipitate  with  barium  chloride.  On  evapo- 
rating the  solution,  the  salt  crystallises  — 

HNa2PO4+  H3PO4  =  2H2NaPO4. 
The  aqueous  solution  of  this  salt  is  acid. 


Lit Jiium 


543 


Hydrogen  Sodium  Ammonium  Phosphate  (microcosmic  salt}, 
HNa(NH4)PO4,4H2O,  is  obtained  by  adding  a  strong  solution  of 
common  sodium  phosphate  to  ammonium  chloride  — 


HNa2PO4  +  NH4Cl  =  NaCl  +  HNa(NH4)PO4. 

The  orthophosphates  are  readily  converted  into  pyro-  and  meta- 
phosphates  (see  page  476). 


LITHIUM. 

Symbol,  Li.     Atomic  weight =7.00. 

Occurrence.— Lithium  is  only  found  in  combination  with  other 
elements.  It  is  a  constituent  of  a  few  somewhat  rare  minerals,  as 
petalite,  30SiO2,4Al2O3,Na2O,2Li2O  ;  spodumene,  15SiO2,4Al2O3> 
3Li2O  ;  lepidolite,  or  lithium  mica,  9SiO2,3Al2O3,K2O,4LiF. 

By  means  of  the  spectroscope,  lithium  compounds  have  been 
detected  in  sea-water,  and  in  most  spring  and  river  waters.  In  a 
few  cases  spring  waters  are  met  with  which  contain  considerable 
quantities  of  lithium  salts.  Thus,  W.  A.  Miller  found  as  much  as 
0.372  gramme  of  lithium  chloride  in  i  litre  of  the  water  of  a  spring 
near  Redruth  in  Cornwall. 

Mode  Of  Formation.— Lithium  is 'obtained  by  the  electrolytic 
decomposition  of  the  fused  chloride.  For  this  purpose  the  dry 
salt  is  heated  in  a  porcelain  crucible,  when  it  melts  at  a  low  red 
heat  to  a  mobile  liquid.  A  rod  of  gas  carbon  is  made  the  positive 
electrode  ;  and  a  stout  iron  wire,  one  end  of  which  is  flattened  out, 
is  used  for  the  negative  pole,  upon  which  the  lithium  is  collected. 
On  passing  an  electric  current  through  the  molten  chloride,  the 
metal  forms  as  a  bright  globule  upon  the  negative  electrode.  The 
wire  is  withdrawn  and  quickly  dipped  beneath  petroleum,  and  the 
solidified  globule  of  lithium  is  then  cut  off  with  a  knife.  The 
reduced  metal,  in  its  passage  from  the  crucible  to  the  petroleum, 
is  protected  from  oxidation  by  the  film  of  fused  chloride  which 
coats  it. 

Properties. — Lithium  is  a  soft,  silver-white  metal,  which  soon 
tarnishes  on  exposure  to  the  air.  It  is  easily  cut  with  a  knife, 
being  softer  than  lead,  but  harder  than  sodium.  It  may  be  pressed 
into  wire,  and  two  pieces  of  the  metal  may  be  made  to  adhere, 
or  welded  together,  at  the  ordinary  temperature.  Lithium  is  the 
lightest  known  solid,  its  specific  gravity  being  0.59.  Its  extreme 


544  Inorganic  Chemistry 

lightness  is  illustrated  by  the  fact  that  the  metal  floats  upon 
petroleum,  a  liquid  which  itself  floats  upon  water.  Lithium  melts 
at  1 80°,  and  at  a  higher  temperature  it  takes  fire  and  burns  with 
a  bright  white  light.  Lithium  decomposes  water  at  the  ordinary 
temperature,  liberating  hydrogen  and  forming  lithium  hydroxide, 
Li  HO  ;  but  when  a  fragment  of  the  metal  is  thrown  upon  cold 
water  it  does  not  melt,  and  even  with  boiling  water  the  action  is 
not  attended  by  inflammation  of  the  hydrogen. 

When  strongly  heated  in  nitrogen  the  two  elements  unite,  with 
feeble  combustion,  forming  lithium  nitride,  NLi3. 

Lithium  Oxide,  Li2O,  is  formed  when  the  metal  burns  in  the 
air.  It  is  also  obtained  by  heating  the  nitrate.  It  dissolves  in 
water,  forming  lithium  hydroxide,  LiHO. 

Lithium  Hydroxide  is  produced  by  the  prolonged  boiling  of 
lithium  carbonate  with  milk  of  lime,  the  carbonate  of  this  metal, 
unlike  potassium  and  sodium  carbonates,  being  only  very  slightly 
soluble  in  water. 

Lithium  Carbonate,  Li2CO3,  is  obtained  as  a  white  precipitate 
when  a  solution  of  either  potassium,  sodium,  or  ammonium  car- 
bonate is  added  to  a  solution  of  either  chloride  or  nitrate  of 
lithium.  The  compound  is  only  slightly  soluble  in  cold  water,  100 
parts  of  water  at  13°  dissolving  0.77  part  of  the  carbonate. 

Lithium  Phosphate,  Li3PO4,  is  precipitated  as  a  crystalline 
powder,  by  the  addition  of  hydrogen  disodium  phosphate  to  a 
solution  of  a  lithium  salt.  In  the  presence  of  sodium  hydroxide  the 
precipitation  is  complete,  and  the  formation  of  this  compound  is 
employed  as  a  quantitative  method  for  estimating  lithium.  The 
crystals  contain  2H2O,  which  they  lose  when  heated.  Lithium 
phosphate  is  soluble  in  nitric,  hydrochloric,  and  phosphoric  acidsi 
and  from  the  latter  solution,  on  evaporation,  the  dihydrogen 
phosphate  is  deposited  (H2LiPO4)  as  deliquescent  and  very  soluble 
crystals.  The  chloride,  nitrate,  and  sulphate  of  lithium  are  obtained 
by  dissolving  the  carbonate  in  the  respective  acids.  The  salts 
are  readily  soluble  in  water. 

Rubidium  and  Caesium.* — These  two  rare  elements,  which  were  first  dis- 
covered by  Bunsen  in  the  waters  of  Diirkheim,  in  the  years  1 860-61,  are  met 
with,  associated  with  sodium  and  potassium,  in  certain  minerals,  such  as 
lepidolites  (lithium  mica),  porphyrites,  and  in  carnallite.  They  are  also  found 

*  For  detailed  descriptions  of  these  elements  and  their  compounds,  the 
student  is  referred  to  larger  works. 


Ammonium  Salts  545 

in  many  mineral  waters,  in  the  mother-liquors  from  the  evaporation  of  sea- 
water,  and  in  the  ashes  of  plants.  Although  widely  distributed,  the  quantities 
present  are  extremely  minute,  one  of  the  richest  lepidolites  in  which  these 
metals  occur  containing  only  0.24  per  cent,  of  rubidium  oxide. 

The  rare  mineral  pollux,  a  silicate  of  aluminium  and  caesium,  containing 
also  iron  calcium  and  sodium,  is  the  only  known  mineral  in  which  either  of 
these  two  elements  occurs  as  an  essential  constituent.  The  analysis  of  Pisani 
(1864)  gives  34.07  per  cent,  of  caesium  oxide  in  this  substance. 

Rubidium  is  obtained  by  heating  the  carbonate  with  carbon  (the  charred 
tartrate),  as  in  the  older  method  for  the  preparation  of  sodium  and  potassium. 

Caesium  cannot  be  isolated  by  this  reaction,  but  is  obtained  by  the  electro- 
lysis of  the  fused  cyanide,  Cs(CN)  (mixed  with  barium  cyanide  in  order  to 
render  it  more  readily  fusible).  Rubidium  melts  at  38.5°,  caesium  at  26.5°. 

Rubidium  gives  a  green  vapour,  and  when  sublimed  in  a  vacuous  tube  yields 
a  thin  film  of  metal,  which  appears  deep  blue  by  transmitted  light  :  when 
slowly  sublimed  in  this  way  the  metal  forms  small  needle-shaped  crystals. 
The  compounds  of  these  metals  closely  resemble  those  of  potassium,  from 
which  they  can  only  be  distinguished  by  the  different  spectra  they  give. 

AMMONIUM  SALTS. 

The  monovalent  group  or  radical  (NH4)  is  capable  of  replacing 
one  atom  of  hydrogen  in  acids,  thereby  giving  rise  to  a  series  of 
salts  which  are  closely  analogous  to,  and  are  isomorphous  with, 
those  of  potassium.  The  radical  (NH4),  to  which  the  name 
ammonium  is  given,  has  never  been  isolated.  United  to  an  atomic 
electric  charge  it  constitutes  the  anion  NH4',  ammonion,  which 
closely  resembles  sodion  and  potassion.  When  an  amalgam  of 
sodium  and  mercury  is  thrown  into  a  solution  of  ammonium 
chloride,  the  mercury  swells  up  into  a  honeycombed  or  sponge- 
like  mass,  which  floats  upon  the  surface  of  the  liquid.  This  so- 
called  ammonium  amalgam  was  at  one  time  thought  to  be  a  true 
amalgam  of  mercury  with  the  metallic  radical  ammonium.  It  is 
now  generally  believed  to  consist  of  mercury  which  is  simply 
inflated  by  the  evolution  of  hydrogen  and  ammonia  gas.  When 
this  sponge-like  substance  is  subjected  to  changes  of  pressure,  it 
is  found  to  contract  and  expand  in  conformity  to  Boyle's  law  :  its 
formation  may  be  represented  by  the  equation  — 


In  the  course  of  a  few  minutes  the  inflated  mass  shrinks  down, 
and  ordinary  mercury  remains  at  the  bottom  of  the  solution, 
hydrogen  and  ammonia  having  been  rapidly  evolved. 

The  ammonium  salts  are  obtained  for  the  most  part  from  the 
ammoniacal  liquor  of  the  gasworks.  The  material  is  treated  with 

2  M 


546  Inorganic  Chemistry 

lime,  and  distilled  ;  and  the  ammonia  so  driven  off  is  absorbed  in 
sulphuric  or  hydrochloric  acid,  giving  rise  to  ammonium  sulphate 
or  chloride. 

Ammonium  Chloride  (sal  ammoniac},  NH4C1.— The  product 
obtained  by  absorbing  ammonia  from  gas  liquor  in  hydrochloric 
acid. is  purified  by  sublimation.      The  crude   material  is  heated" 
in  large  iron  pots,  covered  with  iron  dome-shaped  vessels,  into 


FIG.  141. 

which  the  substance  sublimes.  Ammonium  chloride  crystallises  in 
arborescent  or  fern-like  crystals  (Fig.  141),  consisting  of  groups  of 
small  octahedra  belonging  to  the  regular  system. 

100  parts  of  water  at  10°  dissolve  32.8  parts,  and  at  100°,  77  parts 
of  the  salt.  On  boiling  the  aqueous  solution,  dissociation  to  a 
small  extent  takes  place,  and  a  portion  of  the  ammonia  escapes 
with  the  steam  ;  the  solution  at  the  same  time  becoming  slightly 
acid. 

Ammonium  Sulphate  (NH4)2SO4.— The  product  obtained  by 


Ammonium  Carbonate  547 

the  absorption  of  ammonia  obtained  from  gas  liquors  by  sul- 
phuric acid  is  purified  by  recrystallisation,  when  it  forms  colourless 
rhombic  crystals,  isomorphous  with  potassium  sulphate.  100  parts 
of  water  at  the  ordinary  temperature  dissolve  50  parts  of  the  salt. 
The  chief  use  of  ammonium  sulphate  is  for  agricultural  purposes, 
as  a  manure  ;  and  for  this  use  the  crude  salt,  as  first  obtained, 
which  is  usually  more  or  less  coloured  with  tarry  matters,  is  em- 
ployed. Ammonium  sulphate  is  also  used  for  the  preparation  of 
ammonia  alum  and  other  ammonium  compounds,  as  well  as  in 
'the  ammonia-soda  process. 

Ammonium  Carbonates.  —  Commercial  ammonium  carbonate 
(sal  'volatile}  is  obtained  by  heating  a  mixture  of  ammonium 
sulphate  and  ground  chalk  to  redness  in  horizontal  iron  retorts  or 
cylinders,  and  conducting  the  vapours  into  leaden  receivers  or 
chambers,  where  the  carbonate  condenses  as  a  solid  crust.  It  is 
afterwards  purified  by  resublimation,  when  it  is  obtained  as  a 
white  fibrous  mass.  This  substance  is  a  mixture  of  hydrogen 
ammonium  carbonate,  H(NH4)CO3,  and  ammonium  carbamate, 
(NH4)CO2(NH2),  and  smells  strongly  ammoniacal.  When  treated 
with  alcohol  the  ammonium  carbamate  is  dissolved,  leaving  the 
carbonate  behind. 

Normal  Ammonium  Carbonate,  (NH4)2CO3,  is  obtained  from 
the  commercial  compound,  by  passing  ammonia  gas  into  a  strong 
aqueous  solution,  or  by  digesting  the  compound  in  strong  aqueous 
ammonia.  The  carbamate  present  is  converted  into  normal  car- 
bonate by  the  action  of  the  water,  thus  — 


and  the  ammonia  converts  the  bicarbonate  into  the  normal  salt, 
thus— 

H(NH4)CO3  +  NH3  =  (NH4)2CO3. 

Normal  ammonium  carbonate  on  exposure  to  the  air  gives  off 
ammonia,  and  passes  back  into  hydrogen  ammonium  carbonate. 
When  heated  to  60°  the  salt  breaks  up  into  carbon  dioxide, 
ammonia,  and  water. 

Hydrogen  Ammonium  Carbonate,  H(NH4)CO3,  may  also  be 
obtained  by  passing  carbon  dioxide  into  a  solution  of  the  normal 
salt— 

(NH4)2C03  +  C02  +  H2O  =  2H(NH4)CO3. 

It  forms  large  lustrous  crystals  belonging   to  the  orthorhombic 


548  Inorganic  Chemistry 

system,  which,  when  dry,  do  not  smell  of  ammonia.  100  parts  of 
water  at  15°  dissolve  12.5  parts  of  this  salt.  At  ordinary  tempera- 
tures this  solution  on  exposure  to  the  air  slowly  gives  off  carbon 
dioxide,  and  becomes  alkaline  ;  and  when  heated  above  36°  the 
liquid  begins  to  effervesce,  owing  to  the  rapid  evolution  of  carbon 
dioxide.  This  salt  forms  with  the  normal  carbonate  a  double  salt 
analogous  to  sodium  sesquicarbonate,  and  having  the  composition 
CNH4)2CO3,2H(NH4)CO3,H2O. 

Ammonium  Thioeyanate,  NH4S(CN),  is  prepared  by  adding 
aqueous  ammonia  to  an  alcoholic  solution  of  carbon  disulphide, 
and  allowing  the  mixture  to  stand,  when  ammonium  thiocarbonate 
is  formed,  thus  — 


On  heating  this  solution,  the  ammonium  thiocarbonate  is  de- 
composed with  evolution  of  sulphuretted  hydrogen  — 

(NH4)2CS3=2H2S  +  NH4S(CN). 

Ammonium  thiocyanate  (known  also  as  ammonium  sulpho- 
cyanate)  forms  colourless  crystals,  which  are  extremely  soluble 
both  in  water  and  alcohol.  The  solution  in  water  is  attended  with 
considerable  absorption  of  heat  :  thus,  if  20  grammes  of  the  salt 
be  dissolved  in  25  cubic  centimetres  of  water  at  18°,  the  temperature 
of  the  liquid  falls  to  -  13°. 


CHAPTER   V 
THE  ELEMENTS  OF   GROUP  I.   (FAMILY  £.) 

Copper,  Cu 63.6 

Silver,  Ag 107.88 

Gold,  Au     .......     197.2 

THE  elements  of  this  family  present  many  striking  contrasts  to 
those  of  the  other  family  belonging  to  the  first  group.  These 
three  metals  are  not  acted  upon  by  oxygen,  or  by  water,  at 
ordinary  temperatures  ;  they  are  all  found  native  in  the  un- 
combined  state,  and  on  this  account  are  amongst  the  earliest 
metals  known  to  man.  The  alkali  metals,  on  the  other  hand,  are 
instantly  oxidised  on  exposure  to  air,  they  decompose  water  at 
the  ordinary  temperature,  are  never  found  native,  and  are  amongst 
the  most  recently  discovered  metals.  With  the  exception  of 
sodium  and  potassium,  which  are  used  in  a  few  manufacturing 
processes,  the  alkali  metals,  as  such,  are  of  little  practical  service 
to  mankind,  whilst  the  metals  of  this  family  are  amongst  the  most 
useful  of  all  the  metals,  and  are  the  three  universally  adopted  for 
coinage.  Many  of  the  compounds  of  the  elements  of  this  family 
are  similarly  constituted  to  those  of  the  alkali  metals  :  thus,  with 
oxygen  and  with  sulphur  we  have  Cu2O,  Ag2O,  Au2O,  and  Cu2S, 
Ag2S,  Au2S,  corresponding  to  Li2O  and  K2S. 

With  the  halogens  they  all  form  compounds  of  the  type 
RX.  Although  the  three  elements,  copper,  silver,  and  gold,  fall 
into  the  same  family  upon  the  basis  of  the  periodic  classification 
of  the  elements,  they  are  in  many  respects  widely  dissimilar. 
Thus,  silver  is  consistently  monovalent,  while  copper  is  divalent, 
forming  compounds  of  the  type  CuX2,  and  gold  is  trivalent,  giving 
compounds  AuX3.  The  chlorides,  AgCl  and  Cu2Cl2,  on  the  other 
hand,  are  both  insoluble  in  water,  are  both  soluble  in  ammonia, 
and  both  absorb  ammonia. 

In  many  of  their  physical  attributes,  these  metals  show  a  regular 


550  Inorganic  Chemistry 

gradation  in  their  properties.  Thus,  as  regards  malleability  and 
ductility,  silver  is  intermediate  between  copper  and  gold,  the 
latter  possessing  these  properties  in  the  highest  degree.  With 
respect  to  their  tenacity,  silver  is  again  intermediate,  copper  being 
the  most,  and  gold  the  least  tenacious  of  the  three. 


COPPER. 

Symbol,  Cu.     Atomic  weight =63. 6. 

Occurrence. — Copper  is  found  in  the  elementary  condition  in 
various  parts  of  the  world,  notably  in  the  neighbourhood  of  Lake 
Superior,  where  native  copper  occurs  in  enormous  masses.  In 
combination,  copper  is  a  very  abundant  element,  and  is  widely 
distributed,  the  most  important  of  these  natural  compounds  being 
the  following — 


Ruby  ore    .         .  .     Cu.2O. 

Copper  glance   .  .     Cu2S. 

^  .  (  CuoS.Fe2So  or 

Copper  pyntcs   .  .  {  < 


Purple       copper  (  3Cu2S,FeoS3  or 

ore  .         .         .  (  Cu3FeS3. 
Malachite  .         .     CuCO3,Cu(HO)2. 
Azurite        .         .     2CuCO3,Cu(HO)2. 

Modes  of  Formation. — The  methods  by  which  copper  is 
obtained  from  its  ores  vary  with  the  nature  of  the  ore.  From 
ores  containing  no  sulphur,  such  as  the  carbonates  and  oxidq 
the  metal  may  be  obtained  by  a  method  known  as  the  reducing 
process,  which  consists  in  smelting  down  the  ore  in  a  blast-furnace 
with  coal  or  coke,  when  the  metal  is  reduced  according  to  the 
equation — 

Cu2O  +  C  =  CO  +  2Cu. 

In  the  case  of  mixed  ores,  containing  sulphides,  the  process 
(known  as  the  English  method)  consists  of  six  distinct  stages — 

(i.)  The  ores,  which  contain  on  an  average  30  per  cent,  of  iron 
and  13  of  copper  (the  remainder  being  chiefly  sulphur  and  silica), 
are  first  calcined  ;  usually  in  a  reverberatory  furnace,  whereby  a 
portion  of  the  sulphur  is  burnt  to  sulphur  dioxide,  and  the  metals 
are  partially  oxidised. 

(2.)  The  second  step  consists  in  fusing  the  calcined  ore  ;  when 
the  copper  oxides,  formed  during  calcination,  react  upon  a  portion 
of  the  ferrous  sulphide,  with  the  formation  of  cuprous  sulphide 
and  ferrous  oxide,  thus — 

Cu2O  +  FeS  =  Cu2S  +  FeO. 


Copper  551 

The  oxide  of  iron  combines  with  the  silica  already  present  (or 
which  is  added  in  the  form  of  metal-slag  obtained  from  the  fourth 
step)  to  form  a  fusible  silicate  of  iron,  or  slag,  which  contains 
little  or  no  copper.  This  is  run  off,  and  a  fused  regulus  remains, 
consisting  of  cuprous  and  ferrous  sulphides,  known  as  coarse-metal, 
and  containing  from  30  to  35  per  cent,  of  copper.  This  molten 
regulus,  which  has  a  composition  very  similar  to  copper  pyrites, 
is  usually  allowed  to  flow  into  water,  whereby  it  is  obtained  in  a 
granulated  condition  favourable  for  the  next  operation. 

(3.)  The  third  step  consists  in  calcining  the  granulated  coarse- 
metal  ;  the  result,  as  in  the  first  calcination,  being  the  removal  of 
a  part  of  the  sulphur  as  sulphur  dioxide,  and  the  partial  oxidation 
of  the  metals. 

(4.)  The  calcined  mass  is  next  fused  along  with  refinery-slag, 
which  results  in  the  production  of  a  regulus  consisting  of  nearly 
Dure  cuprous  sulphide,  the  greater  part  of  the  iron  having  passed 
into  the  slag  (known  as  metal-slag).  This  regulus,  called  Jine- 
metal,  or  white-metal,  contains  from  60  to  75  per  cent,  of  copper. 

(5.)  The  fifth  operation  consists  in  roasting  the  "white-metal'* 
in  a  reverberatory  furnace.  A  portion  of  the  cuprous  sulphide  is 
here  oxidised  into  cuprous  oxide,  which,  as  the  temperature  rises, 
reacts  upon  another  portion  of  cuprous  sulphide,  thus  — 


At  the  same  time  any  remaining  ferrous  sulphide  is  converted  into 
oxide,  thus  — 

3Cu2O  +  FeS  =  6Cu  +  FeO  +  SO2. 

The  metallic  copper  so  obtained  presents  a  blistered  appearance, 
and  on  this  account  is  known  as  blister-copper. 

(6.)  This  impure  copper  is  lastly  subjected  to  a  refining  process. 
For  this  purpose  it  is  melted  down  upon  the  hearth  of  a  reverbera- 
tory furnace,  in  an  oxidising  atmosphere.  The  impurities  present 
in  the  metal,  such  as  iron,  lead,  and  arsenic,  are  the  first  to  oxidise  ; 
and  the  oxides  either  volatilise  or  combine  with  the  siliceous  matter 
of  which  the  furnace  bed  is  composed,  forming  a  slag,  which  is 
removed.  The  oxidation  is  continued  until  the  copper  itself  begins 
to  oxidise,  when  the  oxide  so  formed  reacts  upon  any  remaining 
cuprous  sulphide  with  the  reduction  of  copper  and  the  evolution  of 
sulphur  dioxide,  according  to  the  above  equation.  The  metal  at 
this  stage  is  termed  dry  copper  j  and  in  order  to  reduce  the  copper 


552  Inorganic  Chemistry 

oxide  which  it  still  contains,  the  molten  mass  is  stirred  with  poles 
of  wood,  and  a  quantity  of  anthracite  is  thrown  upon  the  surface  to 
complete  the  reducing  process. 

Wet  Process. — Copper  is  extracted  from  the  burnt  pyrites, 
obtained  in  enormous  quantities  in  the  manufacture  of  sulphuric 
acid,  which  contains  about  3  per  cent,  of  copper.  Although  too 
poor  in  copper  to  be  submitted  to  the  smelting  process,  it  is 
found  that  when  calcined  with  12  to  15  per  cent,  of  common  salt, 
the  copper  is  all  converted  into  cupric  chloride.  On  lixiviating  the 
calcined  mass  with  water,  the  cupric  chloride  goes  into  solution,  and 
metallic  copper  can  be  precipitated  from  it  by  means  of  scrap-iron 
or  by  electrolysis. 

Properties. — Copper  is  a  lustrous  metal,  having  a  characteristic 
reddish-brown  colour.  The  peculiar  copper-red  colour  of  the  metal 
is  best  seen  by  causing  the  light  to  be  several  times  reflected  from 
the  surface  before  reaching  the  eye. 

Native  copper  is  occasionally  found  crystallised  in  regular  octa- 
hedra,  and  small  crystals  of  the  same  form  may  be  artificially 
obtained  by  the  slow  deposition  of  the  metal  from  solutions  of  its 
salts  by  processes  of  reduction. 

Copper  is  an  extremely  tough  metal,  and  admits  of  being  drawn 
into  fine  wire,  and  hammered  out  into  thin  leaf.  Its  ductility  and 
malleability  are  greatly  diminished  by  admixture  with  even  minute 
quantities  of  impurities.  When  heated  nearly  to  its  melting-point, 
copper  becomes  sufficiently  brittle  to  be  powdered.  The  specific 
gravity  of  pure  copper,  electrolytically  deposited,  is  8.945,  which 
by  hammering  is  increased  to  8.95. 

Copper  is  only  slowly  acted  upon  by  exposure  to  dry  air 
at  ordinary  temperatures  ;  but  in  the  presence  of  atmospheric 
moisture  and  carbon  dioxide  it  becomes  coated  with  a  greenish 
basic  carbonate.  When  heated  in  air  or  oxygen,  it  is  converted 
into  black  cupric  oxide,  which  flakes  off  the  surface  in  the  form  of 
scales.  When  volatilised  in  the  electric  arc,  copper  gives  a  vapour 
having  a  rich  emerald-green  colour. 

Copper  is  readily  attacked  by  nitric  acid,  either  dilute  or  con- 
centrated, with  the  formation  of  copper  nitrate  and  oxides  of 
nitrogen  (page  246). 

Dilute  hydrochloric  and  sulphuric  acids  are  without  action  upon 
copper  when  air  is  excluded,  but  slowly  attack  it  in  the  presence 
of  air,  or  in  contact  with  platinum.  Cold  concentrated  sulphuric 
acid  does  not  act  upon  copper  ;  but  when  heated,  copper  sulphate 


Cuprous  Oxide  553 

and  sulphur  dioxide  are  formed,  with  the  simultaneous  production 
of  varying  quantities  of  cuprous  and  cupric  sulphides,  which 
remain  as  a  black  residue  (page  416). 

Finely  divided  copper  is  slowly  dissolved  by  boiling  concen- 
trated hydrochloric  acid,  with  evolution  of  hydrogen  and  formation 
of  cuprous  chloride  — 

=  Cu2Cl2  +  H2  or 


In  the  presence  of  air,  copper  is  acted  upon  by  a  solution  of 
ammonia,  the  oxide  dissolving  in  the  ammonia  forming  a  deep 
blue  solution. 

Copper  is  an  extremely  good  electric  conductor,  being  only 
second  to  silver  in  this  respect  ;  it  is  therefore  extensively  em- 
ployed for  cables,  or  leads,  for  purposes  of  telegraphy  and  electric 
lighting. 

Copper  possesses  the  property,  in  a  high  degree,  of  being  de- 
posited in  a  coherent  form  by  the  electrolysis  of  solutions  of  its 
salts.  On  this  account  it  is  extensively  used  in  processes  of 
electrotyping. 

Alloys  of  Copper.  —  The  most  extensive  use  of  copper  is  in 
the  formation  of  certain  alloys,  many  of  which  are  of  great  technical 
value.  The  following  are  among  the  most  important  :  — 

English  brass  .         .         .  Copper  2  parts  Zinc  I  part 

Dutch  brass  (Tombac}  ,,       5      „  „     i    „ 

Muntz  metal    ...          „       3      „  „     I    „ 

Gun  metal        ...          „       9      „  Tin  I    „ 

Aluminium  bronze  .  „       9      „  Aluminium  I    „ 

Oxides  of  Copper.  —  Two  oxides  of  copper  are  well  known, 
namely,  cuprous  oxide  (copper  sub-oxide\  Cu2O,  and  cupric  oxide 
(copper  monoxide).  CuO. 

Cuprous  Oxide,  Cu2O,  occurs  native  as  red  copper  ore.  It  is 
formed  when  finely  divided  copper  is  gently  heated  in  a  current 
of  air,  or  when  a  mixture  of  cuprous  chloride  and  sodium  carbonate 
is  gently  heated  in  a  covered  crucible. 


Cuprous  oxide  is  also  obtained  when  an  alkaline  solution  of  a 
copper  salt  is  reduced  by  grape  sugar. 


554  Inorganic  Chemistry 

Cuprous  oxide  is  insoluble  in  water  ;  it  is  converted  into  cuprous 
chloride  by  strong  hydrochloric  acid.  Nitric  acid  converts  it  into 
cupric  nitrate  with  the  evolution  of  oxides  of  nitrogen.  When  acted 
upon  by  dilute  sulphuric  acid,  it  is  partly  reduced  to  metallic  copper 
and  partly  oxidised  into  copper  sulphate,  thus  — 


When  heated  with  the  strong  acid  it  is  entirely  oxidised,  thus  — 


Cuprous  oxide  fuses  at  a  red  heat,  and  when  melted  with  glass, 
imparts  to  the  latter  a  rich  ruby-red  colour. 

Cupric  Oxide,  CuO,  occurs  as  the  rather  rare  mineral,  tenorite. 
It  is  formed  when  copper  is  strongly  heated  in  the  air  or  in  oxygen, 
or  by  gently  igniting  either  the  nitrate,  carbonate,  or  hydroxide. 
It  is  a  black  powder,  which  rapidly  absorbs  moisture  from  the 
air.  When  heated,  it  first  cakes  together  and  finally  fuses, 
giving  up  a  part  of  its  oxygen,  and  leaving  a  residue  consisting 
of  CuO,2Cu2O. 

When  heated  in  a  stream  of  carbon  monoxide,  marsh  gas,  or 
hydrogen,  it  is  reduced  to  the  metallic  state.  Similarly,  when 
mixed  with  organic  compounds  containing  carbon  and  hydrogen, 
it  oxidises  these  elements  to  carbon  dioxide  and  water,  itself  being 
reduced  :  on  this  property  depends  its  use  in  the  ultimate  analysis 
of  organic  compounds. 

Cuprie  Hydroxide,  Cu(HO)2,  is  the  pale  blue  precipitate  pro- 
duced when  sodium  or  potassium  hydroxide  is  added  in  excess  to  a 
solution  of  a  copper  salt.  The  compound,  when  washed,  may  be 
dried  at  100°  without  parting  with  water  ;  but  if  the  liquid  in  which 
it  is  precipitated  be  boiled,  the  compound  blackens,  and  is  con- 
verted into  a  hydrate  having  the  composition  Cu(HO)2,2CuO. 
Cupric  hydrate  dissolves  in  ammonia,  forming  a  deep  blue  liquid, 
which  possesses  the  property  of  dissolving  cellulose  (cotton  wool, 
filter  paper,  &c.). 

Salts  of  Copper.  —  Copper  forms  two  elementary  ions,  monocu- 
prion  Cu'  and  dicuprion  Cu",  giving  rise  to  two  series  of  salts, 
namely,  cuprous  and  ctipric  salts.  The  former,  which  are  colour- 
less, readily  pass  by  oxidation  into  cupric  salts,  and  serve  therefore 
as  powerful  reducing  agents,  and  are  mostly  insoluble  in  water. 
The  cupric  salts  in  the  hydrated  condition  are  either  blue  or 
green  in  colour;  the  anhydrous  cupric  salts  are  colourless  or 


Cupric  Chloride  555 

yellow.  The  normal  salts  are  mostly  soluble  in  water.  Copper 
salts  impart  to  a  non-luminous  flame  a  blue  or  green  colour,  and 
on  this  account  are  employed  in  pyrotechny.  The  soluble  salts 
are  poisonous. 

Cuprous  Chloride,  Cu.2Cl2,  may  be  obtained  by  dissolving 
cuprous  oxide  in  hydrochloric  acid.  It  is  more  readily  prepared 
by  boiling  a  solution  of  cupric  chloride  in  hydrochloric  acid, 
with  copper  turnings  or  foil.  The  nascent  hydrogen,  liberated  by 
the  action  of  the  hydrochloric  acid  upon  the  copper,  reduces  the 
cupric  chloride  to  cuprous  chloride.  The  liquid  is  then  poured 
into  water,  which  causes  the  precipitation  of  the  cuprous  chloride 
as  a  white  crystalline  powder. 

A  mixture  of  zinc  dust  and  copper  oxide  added  to  strong  hydro- 
chloric acid  also  yields  cuprous  chloride,  the  nascent  hydrogen  in 
this  case  being  derived  from  the  action  of  the  acid  upon  the  zinc, 
and  this  causes  the  reduction  of  cupric  chloride  formed  by  the 
action  of  the  acid  upon  the  cupric  oxide. 

Cuprous  chloride  melts  when  heated,  and  volatilises  without 
decomposition.  It  is  insoluble  in  water,  but  dissolves  in  hydro- 
chloric acid,  ammonia,  and  alkaline  chlorides.  These  solutions,  on 
exposure  to  the  air,  absorb  oxygen,  turning  first  brown,  and  fin- 
ally depositing  a  greenish-blue  precipitate  of  copper  oxychloride, 
CuCl2,3CuO,4H2O.  This  compound  occurs  native  as  the  mineral 
atacamite.  Solutions  of  cuprous  chloride  also  absorb  carbon 
monoxide,  forming  a  crystalline  compound,  believed  to  have  the 
composition,  COCu2Cl2,2H2O.  They  also  absorb  acetylene  (see 
page  318). 

Cuprous  bromide,  Cu2Br2 ;  iodide,  Cu2I2  ;  and  fluoride,  Cu2F2, 
are  also  known. 

Cuprie  Chloride,  CuCl2.  —  This  compound  is  formed  when 
copper  is  dissolved  in  nitro-hydrochloric  acid,  or  when  cupric 
oxide,  carbonate,  or  hydroxide  are  dissolved  in  hydrochloric  acid. 
It  is  also  produced  when  copper  is  burnt  in  chlorine. 

Cupric  chloride  is  readily  soluble  in  water,  forming  a  deep  green 
solution,  which,  on  being  largely  diluted,  turns  blue.  The  salt 
crystallises  in  green  rhombic  prisms,  with  2H2O.  When  heated 
it  loses  its  water,  and  at  a  dull  red  heat  is  converted  into  cuprous 
chloride,  with  evolution  of  chlorine  (see  page  355). 

Cupric  chloride  forms  three  compounds  with  ammonia.  The 
anhydrous  salt  absorbs  ammonia  gas,  forming  a  blue  compound, 
CuCl2,6NH3.  When  ammonia  is  passed  into  aqueous  cupric 


556  Inorganic  Chemistry 

chloride,  the  solution  deposits  deep  blue  crystals  (tetragonal 
pyramids)  of  the  compound,  CuCl2,4NH3,H2O.  Both  these  sub- 
stances, when  moderately  heated,  yield  the  green  compound 
CuCl2,2NH3,  which  at  a  higher  temperature  is  decomposed, 
thus  — 

3Cu2Cl2+6NH4 


Cupric  bromide,  CuBr2,  and  fluoride,  CuF2,  are  known,  but  the 
iodide  is  unknown. 

Cuprie  Nitrate,  Cu(NO3)2,3H2O,  may  be  obtained  by  the 
action  of  nitric  acid  upon  cupric  oxide,  hydroxide,  carbonate,  or 
the  metal  itself.  It  is  deposited  from  the  solution  in  deep  blue 
deliquescent  crystals,  soluble  in  alcohol.  When  heated  to  about 
65°  the  crystals  lose  nitric  acid  and  water,  and  are  converted  into 
the  basic  nitrate,  Cu(NO3)2,3Cu(HO)2.  The  normal  salt,  there- 
fore, cannot  be  obtained  anhydrous.  Cupric  nitrate  is  a  caustic, 
powerfully  oxidising  substance.  If  the  moist  salt  be  rubbed  in  a 
mortar  with  a  quantity  of  tinfoil,  the  tin  is  quickly  converted  into 
oxide,  with  considerable  rise  of  temperature.  When  a  solution 
containing  copper  nitrate  and  ammonium  nitrate  is  evaporated,  the 
mixture  suddenly  deflagrates  when  a  certain  degree  of  concentra- 
tion is  reached. 

Cuprie  Sulphate  (blue  vitriol),  CuSO4,5H2O,  is  the  most 
important  of  all  the  copper  salts.  It  is  formed  when  either  the 
metal  or  the  oxide  is  dissolved  in  sulphuric  acid.  On  a  com- 
mercial scale  it  is  obtained  from  waste  copper  by  first  converting 
the  metal  into  sulphide  by  heating  it  in  a  furnace,  and  throwing 
sulphur  upon  the  red-hot  metal.  Air  is  then  admitted,  and  the 
sulphide  is  thereby  oxidised  into  sulphate,  which  is  dissolved  in 
water  and  crystallised. 

It  is  also  manufactured  from  the  sulphur  ores  of  copper,  by 
roasting  them  under  such  conditions  that  the  iron  is  for  the  most 
part  converted  into  oxide,  while  the  copper  is  oxidised  to  sulphate. 
On  lixiviating  the  roasted  mass  the  copper  sulphate,  with  a  certain 
amount  of  ferrous  sulphate,  is  dissolved  out.  The  ores  may 
also  be  roasted  so  as  to  convert  both  the  metals  into  oxides  ; 
the  mass  is  then  treated  with  "chamber  acid,"  which  dissolves 
copper  oxide,  leaving  the  iron  oxide  for  the  most  part  unacted 
upon. 

Cupric  and  ferrous  sulphates  cannot  be  entirely  separated  by 
crystallisation,  as  a  solution  of  these  salts  deposits  a  double 


Copper  Sulphides  557 

sulphate  of  the  two  metals.  If,  however,  the  amount  of  iron  pre- 
sent is  comparatively  small,  the  first  crop  of  crystals  obtained  is 
moderately  pure  copper  sulphate.  The  copper  is  removed  from 
the  mother-liquors  by  precipitation  upon  plates  of  iron,  and  the 
copper  so  obtained  is  converted  into  sulphide,  as  above  described. 

Copper  sulphate  forms  large  blue  asymmetric  (triclinic)  crystals, 
with  5H2O.  At  100°  it  is  converted  into  a  bluish-white  salt, 
CuSO4,H2O,  and  at  220°  to  240°  it  becomes  anhydrous.  The 
anhydrous  salt  is  white,  and  extremely  hygroscopic,  and  is  used 
both  for  the  detection  and  removal  of  small  quantities  of  water 
in  organic  liquids. 

100  parts  of  water  at  10°  dissolve  36.6  parts,  and  at  100°,  203.3 
parts  of  the  crystallised  salt. 

Several  basic  sulphates  of  copper  are  known  :  thus,  when  the 
normal  salt  is  submitted  to  prolonged  heating,  it  is  converted  into 
an  amorphous  yellow  powder,  consisting  of  CuSO4,CuO,  which 
when  thrown  into  cold  water  forms  an  insoluble  green  compound, 
CuSO4,3Cu(HO)2,  and  on  treatment  with  boiling  water  yields 
CuSO4,2Cu(HO)2.  Copper  sulphate  forms  several  compounds 
with  ammonia.  Thus,  the  anhydrous  salt  readily  absorbs  ammonia 
gas,  forming  the  compound,  CuSO4,5NH3.  When  excess  of 
ammonia  is  added  to  a  solution  of  copper  sulphate,  the  deep  blue 
solution  deposits  blue  crystals  of  CuSO4,H2O,4NH3.  At  150° 
this  compound  is  converted  into  CuSO4,2NH3,  and  at  200°  it  loses 
one  more  molecule  of  ammonia,  leaving  CuSO4,NH3. 

Cuprie  Carbonates. — The  normal  carbonate  has  not  been 
obtained.  The  two  most  important  basic  carbonates  are  (i) 
CuCO3,Cu(HO)2,  occurring  native  as  malachite,  and  obtained  when 
sodium  carbonate  is  added  to  a  solution  of  copper  sulphate  (the 
green  deposit  which  appears  upon  copper  when  exposed  to  atmos- 
pheric moisture  and  carbon  dioxide  (verdigris)  is  the  same  com- 
pound) ;  and  (2)  2CuCO3,Cu(HO)2,  occurring  as  the  mineral  azunte. 

Sulphides  of  Copper. — Two  sulphides  are  known,  correspond- 
ing to  the  two  oxides. 

Cuprous  Sulphide,  Cu2S,  occurs  in  nature  as  copper  glance, 
in  the  form  of  grey  metallic-looking  rhombic  crystals.  It  is  pro- 
duced when  copper  burns  in  sulphur  vapour,  or  when  an  excess 
of  copper  filings  is  heated  with  sulphur. 

Cuprie  Sulphide,  CuS,  is  met  with  in  nature  as  the  mineral  indigo- 
copper.  It  is  obtained  when  either  copper  or  cuprous  sulphide  is 
heated  with  sulphur  to  a  temperature  not  beyond  1 14° ;  so  obtained, 


558  Inorganic  Chemistry 

the  compound  is  blue.     As  a  black  precipitate,  it  is  formed  when 
sulphuretted  hydrogen  is  passed  into  solutions  of  cupric  salts. 


SILVER. 

Symbol,  Ag.     Atomic  weight  — 107.88. 

.Occurrence. — Silver  is  found  uncombined,  occasionally  in 
masses  weighing  several  cwts.  Such  native  silver  usually  contains 
copper,  gold,  and  other  metals. 

Amongst  the  more  important  natural  compounds  of  silver  are 
the  following  : — 

Argentite,  or  silver  glance        .         .  Ag2S. 

Pyrargyrite,  or  ruby  silver  ore         .  3Ag2S,Sb2S3,  or  Ag3SbS3. 

Proustite,  or  light  red  silver  ore      .  3Ag2S,As2S3    „  Ag3AsS3. 

Stephanite 5Ag2S,Sb2S3    „  Ag5SbS4. 

Polybasite 9(Ag2S,Cu2S),Sb2S3,As2S3. 

Stromeyerite    .         ...         .         .  Ag2S,Cu2S. 

Horn  silver      .  ...  AgCl. 

Silver  is  present  also  in  most  ores  of  lead,  notably  with  galena 
(lead  sulphide)  ;  argentiferous  lead  ores  constituting  one  of  the 
main  supplies  of  silver. 

Modes  of  Formation. — This  element  may  be  obtained  from 
its  salts  by  the  electrolysis  of  their  aqueous  solutions.  The  metal 
is  so  readily  reduced  from  its  compounds,  that  many  organic 
substances,  such  as  grape  sugar,  aldehyde,  certain  tartrates,  &c., 
are  capable  of  effecting  its  deposition.  When  a  strip  of  zinc  is 
introduced  into  silver  nitrate  solution,  the  silver  is  at  once  de- 
posited upon  the  zinc  as  a  crystalline  mass. 

Pure  silver  for  analytical  purposes  may  be  prepared  by  pre- 
cipitating silver  chloride,  by  the  addition  of  hydrochloric  acid  to 
a  solution  of  the  nitrate,  and  reducing  the  chloride  by  boiling  with 
sodium  hydroxide  and  sugar,  or  by  means  of  metallic  zinc.  In 
this  way  the  metal  is  obtained  as  a  fine  grey  powder.  The 
chloride  may  also  be  reduced  by  fusion  with  sodium  carbonate, 
when  the  silver  is  obtained  as  a  button  at  the  bottom  of  the 
crucible.  The  methods  by  which  silver  is  obtained  from  its  ores 
are  very  varied  ;  they  may,  however,  be  classed,  under  three  heads, 
namely — 


Silver  559 

1.  Processes   involving   the   use   of  mercury.      (Amalgamation 
processes.) 

2.  Processes  by  means  of  lead. 

3.  Wet  processes. 

(i.)  Amalgamation  Processes.— These  depend  upon  the  fact 
that  certain  compounds  of  silver  are  reduced  by  mercury.  The 
reduced  silver  then  dissolves  in  the  mercury,  forming  an  amalgam, 
from  which  the  silver  is  obtained,  and  the  mercury  recovered  by 
distillation.  The  process,  as  still  carried  on  in  Mexico  and  South 
America,  is  the  following.  The  ore  is  first  crushed  and  then 
ground  to  a  fine  powder  with  water,  and  the  mud  so  obtained  is 
mixed  with  3  to  5  per  cent,  of  common  salt,  and  spread  upon  the 
floor  of  a  circular  paved  space,  the  mixing  being  effected  by  the 
treading  of  mules.  After  the  lapse  of  a  day,  mercury  is  added, 
together  with  a  quantity  of  roasted  pyrites  (known  as  magistral, 
and  consisting  of  a  crude  mixture  of  cupric  and  ferric  sulphates 
and  oxides),  and  the  materials  thoroughly  incorporated.  Fresh 
mercury  is  added  from  time  to  time,  during  the  several  days 
required  for  the  completion  of  the  chemical  decompositions  that 
take  place.  The  exact  nature  of  these  changes  is  not  thoroughly 
understood,  but  it  is  probable  that  they  involve  first  the  formation 
of  copper  chlorides,  by  double  decomposition  between  the  copper 
sulphate  and  sodium  chloride,  and  the  subsequent  action  of  these 
upon  the  silver  sulphide  present  in  the  ore,  thus — 

2CuCl2  +  Ag2S  =  2AgCl  +  Cu2Cl2  +  S. 
Cu2Cl2  +  Ag2S  =  2AgCl  +  Cu2S. 

The  silver  chloride  dissolves  in  the  sodium  chloride  present,  and 
is  reduced  by  the  mercury,  with  the  production  of  mercurous 
chloride  (calomel},  which  is  ultimately  lost  in  the  washing — 

2AgCl  +  2Hg=Hg2Cl2  +  2Ag. 

The  amalgam  is  first  washed,  and  freed  from  adhering  particles 
of  mineral,  and  is  then  filtered  through  canvas  bags,  whereby  the 
excess  of  mercury  is  removed.  The  solid  residue,  containing  the 
silver,  is  then  submitted  to  distillation. 

In  other  amalgamation  processes  the  ore  is  first  roasted  with 
salt,  in  order  to  convert  the  silver  into  chloride.  The  roasted 
ore  is  reduced  to  fine  powder  with  water,  and  introduced  into 


560  Inorganic  Chemistry 

revolving  casks  along  with  scrap  iron,  when  the  chloride  is  reduced 
according  to  the  equation  — 

2AgCl  +  Fe  =  2  Ag  +  FeCl2, 

and  the  reduced  silver  is  then  extracted  by  the  addition  of  mercury, 
with  which  it  amalgamates. 

In  the  modern  amalgamation  process  the  finely  crushed  ore,  with 
water,  is  placed  in  iron  pans  provided  with  revolving  machinery, 
which  serves  the  purpose  of  further  grinding,  and  also  of  mixing. 
When  the  ore  is  reduced  to  an  almost  impalpable  powder,  mercury 
is  added,  and  the  machinery  is  kept  in  operation  for  a  few  hours, 
when  the  amalgamation  is  complete  ;  sometimes  common  salt  and 
copper  sulphate  are  added,  either  together  or  singly.  Their  pre- 
sence does  not  appear  to  be  necessary  to  the  process,  except  in  so 
far  as  they  aid  in  keeping  the  surface  of  the  mercury  clean,  or 
"  quick"  ;  for  in  the  extremely  finely  divided  condition  to  which  the 
ore  is  reduced  in  this  "pan"  amalgamation  process,  the  silver 
sulphide  is  readily  acted  upon  by  mercury,  with  the  formation  of 
mercuric  sulphide  — 


and  the  silver  so  reduced  dissolves  in  the  excess  of  mercury,  from 
which  it  is  finally  separated  by  distillation. 

(2.)  Processes  by  Means  of  Lead.  —  When  silver  ores  are 

smelted  with  lead,  or  with  materials  which  yield  metallic  lead  ;  in 
other  words,  when  silver  ores  are  smelted  with  lead  ores,  an  alloy  of 
silver  and  lead  is  obtained,  from  which  the  silver  can  be  separated. 
When  the  argentiferous  lead  is  rich  in  silver,  the  alloy  is  submitted 
to  cupellation,  which  consists  in  heating  the  metal  in  a  reverbera- 
tory  furnace,  the  hearth  of  which  consists  of  a  movable,  oval-shaped, 
shallow  dish,  made  of  bone  ash,  known  as  a  cupel,  or  test.  The 
alloy  is  fed  into  this  cupel  from  a  melting-pot,  and  a  blast  of  air  is 
projected  upon  the  surface  of  the  molten  metal.  The  lead  is  thus 
converted  into  litharge,  and  the  melted  oxide,  by  the  force  of  the 
blast,  is  made  to  overflow  into  iron  pots.  As  the  oxidation  of  the 
lead  reaches  completion,  the  thin  film  of  litharge  begins  to  exhibit 
iridescent  interference  colours,  which  presently  disappear,  leaving 
the  brilliant  surface  of  the  melted  silver.  The  sudden  appearance 
of  the  bright  metallic  surface  is  known  as  the  flashing  of  silver. 
In  the  case  of  argentiferous  lead  too  poor  in  silver  to  be  directly 


Silver  561 

cupelled,  the  alloy  is  submitted  to  one  of  two  processes  of  con- 
centration, namely,  the  Pattinson  process,  or  the  Parkes's  process. 

The  Pattinson  process  for  desilverising  lead  depends  upon  the 
fact  that  alloys  of  silver  and  lead  have  a  lower  melting-point  than 
pure  lead,  and  therefore  when  argentiferous  lead  is  melted  and 
allowed  to  cool,  the  crystals  which  first  form  consist  of  lead  which 
is  nearly  or  quite  pure,  and  the  greater  part  of  the  silver  is  in 
the  still  liquid  portion.  The  operation  is  carried  out  in  a  row  of 
iron  pots.  A  quantity  of  the  metal  is  melted  in  one  pot,  and  as 
it  cools  the  crystals  which  begin  to  form  are  removed  by  means 
of  a  perforated  iron  ladle  and  transferred  to  the  next  pot  on 
one  side.  This  operation  is  continued  until  a  definite  proportion 
(either  two-thirds  or  seven-eighths,  depending  upon  the  propor- 
tion of  silver)  has  been  removed.  The  residue  is  then  transferred 
to  the  neighbouring  pot  on  the  opposite  side,  and  a  second  charge 
melted  up  in  the  first  pot.  As  the  neighbouring  pots  fill  up  they 
are  similarly  treated,  and  in  this  way  an  alloy,  gradually  becoming 
richer  and  richer  in  silver,  is  passed  along  in  one  direction,  and 
purer  and  purer  lead  is  sent  in  the  opposite  way.  The  rich  alloy 
is  then  cupelled. 

The  Parkers  process  depends  upon  the  fact  that  when  zinc  is 
added  to  a  melted  alloy  of  lead  and  silver,  the  zinc  deprives  the 
lead  of  the  silver,  and  itself  forms  an  alloy  with  it.  The  alloy  of 
zinc  and  silver  rises  to  the  surface  and  is  the  first  portion  to  solidify, 
and  can  be  removed.  The  operation  is  carried  out  in  iron  pots. 
The  argentiferous  lead  is  melted  and  a  quantity  of  zinc  is 
thoroughly  stirred  into  the  molten  mass,  the  amount  of  zinc 
depending  upon  the  richness  of  the  lead.  As  the  mixture  cools, 
the  first  portions  to  solidify  are  skimmed  off  with  a  ladle  and 
transferred  to  another  pot.  These  skimmings,  consisting  of  zinc, 
silver,  and  lead,  are  first  liquated  ;  that  is,  carefully  heated  to  such 
a  temperature  that  the  adhering  lead  melts  and  flows  away  from 
the  less  fusible  zinc  silver  alloy.  The  solid  alloy  is  then  distilled, 
and  the  residue,  consisting  of  silver  and  lead,  is  submitted  to 
cupellation. 

(3.)  Wet  Processes  {Ziervogel  process). — When  argentiferous 
pyrites,  or  an  artificially  formed  regulus  containing  sulphides  of 
silver,  copper,  and  iron  is  roasted,  the  sulphides  are  first  converted 
into  sulphates  ;  and,  as  the  roasting  continues,  first  the  iron,  then 
the  copper,  and  lastly  the  silver  sulphate,  is  converted  into  oxide. 
By  careful  regulation  the  process  is  continued  until  the  whole 

2  N 


562  Inorganic  Chemistry 

of  the  iron  and  a  part  of  the  copper  sulphates  are  decomposed 
On  lixiviating  the  roasted  mass  with  water,  the  silver  sulphate, 
together  with  the  remaining  copper  sulphate,  dissolves.  From 
this  solution  the  silver  is  precipitated  by  scrap  copper. 

The  copper  is  recovered  from  the  solution  by  precipitation  with 
iron. 

The  Percy-Patera  Process. — In  this  method  the  ore  is  roasted 
with  salt  and  the  silver  chloride  so  formed  is  then  extracted  by 
means  of  sodium  thiosulphate — 

Na2S2O3  +  AgCl  =  NaCl  +  NaAgS2O3. 

To  the  solution  so  obtained  sodium  or  calcium  sulphide  is  added, 
which  precipitates  silver  sulphide — 

2NaAgS2O3  +  Na2S  -  Ag2S  +  2Na2S2O3. 

The  silver  sulphide  is  then  reduced  by  being  roasted  in  a  rever- 

beratory  furnace. 

Properties.— Silver  is  a  lustrous 
white  metal  which  appears  yellow 
when  the  light  is  reflected  many 
times  from  its  surface  before  reach- 
ing the  eye.  It  is  unacted  upon  by 
atmospheric  oxygen,  but  quickly 
becomes  tarnished  by  traces  of 
sulphuretted  hydrogen  in  the  air. 
FlG  Silver  has  the  highest  conductivity 

for  heat  and  electricity  of  all  the 

metals.  It  is  extremely  malleable  and  ductile,  being  second  only 
to  gold.  Thin  films  of  silver  appear  blue  by  transmitted  light. 
Silver  melts  at  about  1000°,  and  when  heated  by  the  oxyhydrogen 
flame  may  be  readily  made  to  boil  and  distil.  The  pure  metal 
employed  by  Stas  for  the  determination  of  the  atomic  weight  was 
obtained  by  distillation  in  this  way.  When  volatilised  in  the 
electric  arc,  the  vapour  of  silver  has  a  brilliant  green  colour. 
Molten  silver  absorbs  as  much  as  twenty-two  times  its  volume  of 
oxygen,  which  it  gives  up  again  (with  the  exception  of  0.7  volume) 
on  solidification.  As  the  mass  cools,  the  oxygen  evolved  often 
bursts  through  the  outer  crust  of  solidified  metal  with  consider- 
able violence,  ejecting  portions  of  the  still  liquid  silver  as  irregular 
excrescences,  as  seen  in  Fig.  142.  This  phenomenon  is  known 


Silver  Oxides  563 

as  the  "  spitting  "  of  silver.  Small  quantities  of  admixed  metals 
prevent  the  absorption  of  oxygen. 

Silver  is  readily  soluble  in  nitric  acid,  forming  argentic  nitrate,  with 
liberation  of  oxides  of  nitrogen.  Hot  concentrated  sulphuric  acid 
converts  it  into  argentic  sulphate,  with  formation  of  sulphur  dioxide 
(the  reactions  in  both  cases  being  similar  to  those  with  copper). 

Silver  Alloys. — Silver,  alloyed  with  copper,  is  largely  employed 
for  coinage  and  for  ornamental  purposes.  English  standard  silvef 
contains  925  parts  of  silver  per  1000.  It  is  said,  therefore,  to 
have  a  fineness  of  925.  In  France  three  standards  are  used. 
That  for  coinage  contains  900  parts  per  1000.  For  medals  and 
plate  the  silver  has  a  fineness  of  950,  while  for  jewellery  it  con- 
tains only  800  parts  per  1000. 

Silver-plating". — For  purposes  of  electro-plating,  a  solution  of  silver  cyanide 
in  potassium  cyanide  is  used.  When  a  feeble  electric  current  is  passed 
through  this  solution  (the  article  to  be  silvered  being  the  negative  electrode, 
and  a  plate  of  silver  the  positive),  silver  in  a  coherent  form  is  precipitated 
upon  the  negative  electrode,  thereby  coating  the  object ;  and  cyanogen  is  dis- 
engaged at  the  positive  pole,  where  it  dissolves  the  electrode,  reforming  silver 
cyanide. 

Silver  is  reduced  from  solutions  and  deposited  as  a  coherent  film  by  a 
variety  of  organic  compounds  ;  and  various  methods  based  upon  this  property 
are  in  use  for  obtaining  mirrors  and  silvered  glass  specula  for  optical  pur- 
poses. One  such  method  is  the  following.  Two  solutions  are  prepared, 
thus— 

(t.)  Ten  grammes  of  silver  nitrate  are  dissolved  in  a  small  quantity  of 
water,  and  ammonia  added  until  the  precipitate  dissolves.  The  liquid  is  then 
filtered  and  diluted  up  to  one  litre. 

(2. )  Two  grammes  of  silver  nitrate  are  dissolved  in  a  litre  of  boiling  water, 
and  1.66  gramme  of  Rochelle  salt  (sodium  potassium  tartrate,  NaKC4H4O6) 
are  added  and  the  liquid  filtered.  Equal  volumes  of  these  two  solutions  are 
poured  into  a  shallow  dish,  and  the  glass  to  be  silvered  (after  being  perfectly 
cleaned)  is  laid  in  the  solution.  In  about  twenty  minutes  the  silver  will  have 
formed  a  brilliant  mirror  upon  the  glass.* 

Oxides  of  Silver. — Three  oxides  are  believed  to  exist,  namely — 

Silver  monoxide  ....  Ag2O. 
Silver  peroxide  ....  Ag2O2 
Silver  suboxide  ....  Ag4O  ? 

*  By  the  reduction  of  silver  solutions  in  the  presence  of  certain  organic 
compounds,  Carey  Lea  has  obtained  the  metal  in  the  form  of  a  dark  bronze 
powder,  which,  when  dry,  resembles  burnished  gold.  He  has  also  obtained 
it  exhibiting  bluish-green  and  ruby-red  colours.  The  material  differs  in 
many  of  its  properties  from  ordinary  silver,  and  is  regarded  by  its  discoverer 
as  an  allotropic  form  of  silver  (American  Journal  of  Science,  1891). 


564  Inorganic  Chemistry 

Silver  Monoxide  (argentic  oxide\  Ag2O,  is  obtained  by  adding 
sodium  or  potassium  hydroxide  to  a  solution  of  silver  nitrate.  A 
brown  precipitate  consisting  of  hydrated  oxide  is  obtained  which, 
when  heated,  is  converted  into  the  anhydrous  compound.  It  is 
also  formed  when  silver  chloride  is  boiled  with  a  strong  solution  of 
potassium  hydroxide  — 


Silver  oxide  is  a  black  amorphous  powder  which,  when  heated 
to  260°,  begins  to  give  off  oxygen,  and  become  reduced  to  metallic 
silver.  It  is  a  powerful  oxidising  substance,  and  when  rubbed 
with  sulphur,  red  phosphorus,  sulphides  of  antimony  or  arsenic,  or 
other  readily  oxidised  substances,  it  causes  them  to  ignite. 

Silver  oxide,  although  only  very  slightly  soluble  in  water  (i  part  in 
about  3000),  imparts  to  the  solution  a  distinct  metallic  taste  and  an 
alkaline  reaction. 

It  is  reduced  by  hydrogen  at  100°,  with  formation  of  water  and 
metallic  silver  ;  and  when  brought  into  contact  with  peroxide  of 
hydrogen,  oxygen  is  evolved  and  metallic  silver  formed  (see  p.  227). 

Silver  oxide  is  soluble  in  strong  ammonia,  and,  on  standing,  the 
solution  deposits  black  shining  crystals  of  the  so-called  fulminating 
silver.  When  dry  this  compound  is  extremely  explosive,  and  it 
often  explodes  when  wet.  Fulminating  silver  is  believed  to  be  the 
nitride,  with  the  composition  NAg3. 

Silver  Peroxide,  Ag2O2.  —  When  a  solution  of  silver  nitrate  is  submitted  to 
electrolysis,  a  black  powder,  consisting  of  small  octahedral  crystals,  is  deposited 
upon  the  positive  electrode.  The  same  compound  is  obtained  when  a  plate  of 
silver  is  made  the  positive  electrode  in  the  electrolysis  of  acidulated  water, 
and  also  when  silver  is  acted  upon  by  ozone. 

It  readily  parts  with  oxygen,  and  is  a  still  more  powerful  oxidising  agent 
than  the  monoxide.  It  dissolves  in  aqueous  ammonia  with  the  evolution  of 
nitrogen  — 

3Ag202  +  2NH3=3Ag20  +  3H2O  +  N2. 

Silver  Suboxide,  Ag4O(?).  —  The  black  powder,  obtained  when  silver  citrate 
is  reduced  in  a  current  of  hydrogen  at  100°,  and  potassium  hydroxide  is 
added  to  the  aqueous  solution  of  the  residue,  is  believed  to  have  the  composition 


Silver  Chloride,  AgCl,  is  obtained  as  a  white,  bulky,  curdy 
precipitate  when  a  soluble  chloride  is  added  to  silver  nitrate.  It 
melts  at  451°  to  a  yellowish  liquid,  which,  on  cooling,  congeals  to  a 


Silver  Fluoride  565 

tough  horny  mass  (hence  the  name  horn  silver^  as  applied  to  the 
native  silver  chloride).  The  precipitated  chloride  is  soluble  to  a 
slight  extent  in  strong  hydrochloric  acid,  but  readily  soluble  in 
alkaline  chlorides,  in  ammonia,  and  in  sodium  thiosulphate.  Potas- 
sium cyanide  converts  silver  chloride  into  silver  cyanide,  which 
dissolves  in  the  excess  of  alkaline  cyanide,  forming  the  double 
cyanide  KCN,AgCN.  When  exposed  to  the  light,  silver  chloride 
darkens  in  colour,  assuming  first  a  violet  tint,  and  finally  becoming 
dark  brown  or  black  (see  Photo-salts,  p.  566). 

Silver  chloride  absorbs  large  volumes  of  ammonia,  forming  the 
compound  2AgCl,3NH3  (see  p.  275). 

Silver  Bromide,  AgBr,  is  prepared  similarly  to  the  chloride, 
the  precipitated  compound  having  a  pale  yellow  colour.  It  is  less 
soluble  in  ammonia  than  silver  chloride  ;  in  dilute  ammonia  it  is 
nearly  insoluble.  Silver  bromide  is  decomposed  by  chlorine,  and 
at  a  temperature  of  100°  by  hydrochloric  acid.  At  ordinary  tem- 
peratures this  reaction  is  reversed,  hydrobromic  acid  converting 
silver  chloride  into  the  bromide. 

Dry  silver  bromide  does  not  absorb  gaseous  ammonia.  Silver 
bromide  is  extremely  sensitive  to  the  action  of  light,  and  is  the 
chief  silver  compound  used  in  dry-plate  photography. 

Silver  Iodide,  Agl,  may  be  obtained  by  precipitation  from  silver 
nitrate,  with  a  soluble  iodide  ;  or  by  dissolving  silver  in  strong 
hydriodic  acid.  As  obtained  by  precipitation  it  is  an  amorphous 
yellow  substance,  less  soluble  in  ammonia  than  either  the  bromide 
or  chloride.  It  dissolves  in  hot  hydriodic  acid,  which  on  cooling 
deposits  colourless  crystals  of  Agl, HI  ;  the  addition  of  water  to 
the  solution  precipitates  the  normal  iodide,  Agl.  Silver  iodide 
absorbs  gaseous  ammonia,  forming  a  white  compound,  2AgI,NH3, 
which,  on  free  exposure  to  the  air,  evolves  ammonia,  and  is  recon- 
verted into  the  yellow  iodide. 

Silver  iodide  is  the  most  stable  of  the  three  halogen  compounds. 
When  either  the  chloride  or  bromide  is  treated  with  hydriodic  acid 
or  potassium  iodide,  iodine  replaces  the  other  halogens,  forming 
silver  iodide. 

Silver  Fluoride,  AgF. — This  compound  is  markedly  different 
in  many  respects  from  the  other  halogen  silver  salts.  It  is  obtained 
by  dissolving  silver  oxide  or  carbonate  in  hydrofluoric  acid,  and 
is  deposited  from  the  solution  in  colourless,  tetragonal  pyramids, 
AgF,H2O,  or  in  prisms,  AgF,2H2O.  The  salt  is  extremely  deli- 
quescent, and  very  soluble  in  water.  When  dried  in  vacuo,  the 


566  Inorganic  Chemistry 

salt  AgF,H2O  undergoes  partial  decomposition,  leaving  a  brownish 
residue.  When  heated,  it  is  partially  decomposed,  according  to 
the  equation — 

2AgF,H2O  =  2Ag  +  2HF  +  H2O  +  O. 

The  dry  salt  absorbs  gaseous  ammonia  in  large  quantities,  more 
than  800  times  its  own  volume  being  taken  up  by  the  powdered 
substance. 

Silver  Nitrate,  AgNO3,  is  obtained  by  dissolving  silver  in 
nitric  acid.  It  forms  large  colourless  rhombic  tables,  which  melt 
at  218°,  and  resolidify  to  a  white,  fibrous,  crystalline  mass,  known 
as  lunar  caustic.  Below  a  red  heat  it  gives  off  oxygen,  and  forms 
silver  nitrite  ;  and  at  higher  temperatures  it  is  decomposed  into 
metallic  silver  and  oxides  of  nitrogen.  100  parts  of  water  at  o° 
dissolve  121.9  parts,  and  at  100°,  mo  parts  of  the  crystallised 
salt  ;  the  solution  is  neutral.  In  contact  with  organic  matter, 
silver  nitrate  is  blackened  on  exposure  to  light.  Thus,  when  the 
skin  is  touched  with  a  solution  of  this  salt,  a  few  seconds'  exposure 
to  light  causes  a  brown  or  black  stain.  Owing  to  this  property, 
silver  nitrate  is  employed  for  marking-inks.  Silver  nitrate  absorbs 
gaseous  ammonia,  forming  the  compound  AgNO3,3NH3,  the  ab- 
sorption being  accompanied  with  considerable  rise  of  temperature. 
The  compound  AgNO3,2NH3  is  deposited  as  rhombic  prisms  when 
aqueous  silver  nitrate  is  saturated  with  ammonia. 

Silver  Sulphate,  Ag2SO4,  is  formed  when  silver,  silver  carbo- 
nate, or  silver  oxide  is  dissolved  in  sulphuric  acid.  It  crystallises 
in  rhombic  prisms,  isomorphous  with  sodium  sulphate.  With 
aluminium  sulphate  it  forms  an  alum,  in  which  the  monovalent 
element  silver  takes  the  place  of  potassium  in  common  alum, 
Ag2S04,Al2(S04)3,24H20. 

Photo-salts. — This  name  has  been  applied  by  Carey  Lea  to  the  coloured 
compounds  formed  by  the  action  of  light  upon  silver  chloride,  bromide,  and 
iodide.  The  exact  composition  of  the  compounds  that  are  formed  when  these 
silver  salts  are  exposed  to  light  is  not  definitely  known.  The  change  which  they 
undergo  has  been  attributed  (i)  to  the  partial  reduction  to  metallic  silver ; 
(2)  to  the  formation  of  sub-salts,  such  as  Ag2Cl,  Ag2Br,  with  elimination  of 
chlorine  or  bromine  ;  (3)  to  the  formation  of  oxychloride  or  oxybromide ; 
(4)  to  the  production  of  double  compounds  of  variable  composition  of  the 
sub-salt  with  the  normal  salt. 


Gold  567 


GOLD. 

Symbol,  Au.     Atomic  weight  =  197. 2. 

Occurrence. — This  element  occurs  in  nature  almost  exclusively 
in  the  uncombined  condition,  chiefly  in  quartz  veins  and  in  alluvial 
deposits  formed  by  the  disintegration  of  auriferous  rocks.  It  is 
present  in  small  quantities  in  many  specimens  of  iron  pyrites, 
copper  pyrites,  and  many  lead  ores,  from  which  it  is  often 
profitably  extracted. 

Gold  is  also  met  with  in  the  form  of  an  amalgam,  and  in  com- 
bination with  the  element  tellurium  in  the  minerals  petzite^ 
(AgAu)2Te,  and  sylvanitt,  (AgAu)Te2. 

Extraction. — Gold  is  extracted  from  auriferous  quartz  by  caus- 
ing the  finely-crushed  substance  to  flow,  by  means  of  a  stream 
of  water,  over  amalgamated  copper  plates.  The  gold  particles 
adhere  to  the  mercury,  with  which  they  amalgamate,  and  the 
amalgam  so  obtained  is  carefully  removed  and  distilled. 

From  alluvial  deposits,  the  native  gold  is  separated  by  me- 
chanical washing. 

Gold  is  extracted  from  auriferous  pyrites  by  means  of  chlorine. 
The  ore  is  first  carefully  roasted,  and,  after  being  wetted,  is  exposed 
to  the  action  of  chlorine  gas.  The  gold  is  thereby  converted  into 
the  soluble  auric  chloride,  AuCl3,  which  is  extracted  by  lixiviation, 
and  precipitated  by  the  addition  of  ferrous  sulphate — 

2AuCl3  +  6FeSO4  =  2Au  +  Fe2Cl6  +  2Fe2(SO4)3. 

Native  gold  usually  contains  silver,  from  which  it  may  be  sepa- 
rated by  passing  chlorine  over  the  molten  metal,  in  crucibles  glazed 
with  borax.  The  fused  chloride  of  silver  rises  to  the  surface,  and 
is  prevented  from  volatilising  by  a  covering  of  melted  borax. 
When  the  operation  is  complete,  the  crucible  is  allowed  to  cool, 
when  the  gold  solidifies,  and  the  still  liquid  silver  chloride  is 
poured  off. 

The  Cyanide  Process. — Increasing  quantities  of  gold  are  at 
the  present  time  extracted  by  solution  in  potassium  cyanide.  The 
method  is  specially  advantageous  in  cases  where  the  gold  is  present 
in  the  ore  in  a  very  finely  divided  condition,  and  it  also  possesses 
the  advantage  over  the  "  chlorination  process,"  that  the  preliminary 
operation  of  roasting  is  obviated.  The  crushed  ore  is  treated  with 
a  dilute  solution  of  potassium  cyanide  (containing  from  0.25  to 


568  Inorganic  Chemistry 

I  per  cent,  of  potassium  cyanide),  with  free  exposure  to  the  atmos- 
phere, since  it  has  been  shown  that  atmospheric  oxygen  takes  a 
necessary  part  in  the  action.  The  gold  is  dissolved  in  the  form  of 
a  double  cyanide,  according  to  the  equation  — 


From  this  solution  the  gold  is  precipitated  either  by  means  of 
metallic  zinc  (usually  in  the  form  of  fine  turnings)  or  by  electro- 
lytic deposition.  The  precipitation  by  means  of  zinc  takes  place 
according  to  the  equation  — 

2K  AuCy2  +  Zn  =  K2ZnCy4  +  2Au. 

The  deposit,  after  being  freed  as  far  as  possible  from  zinc,  is 
melted  down  with  a  suitable  flux,  and  yields  an  alloy  containing 
70  to  80  per  cent,  of  gold. 

When  the  gold  is  precipitated  electrolytically,  the  anodes  em- 
ployed are  of  lead  foil.  These  are  finally  melted  down  and  cupelled, 
yielding  gold  of  a  high  degree  of  purity. 

Properties.  —  Gold  is  a  soft  yellow  metal,  which,  when  seen  by 
light  many  times  reflected  from  its  surface,  appears  red.  It  is  not 
acted  upon  by  air  or  oxygen  at  any  temperature,  and  does  not 
decompose  steam.  No  single  acid  is  capable  of  attacking  it 
(except  selenic  acid)  ;  but  it  is  dissolved  by  aqua-regia,  with  for- 
mation of  auric  chloride.  Gold  is  the  most  malleable  and  ductile 
of  all  the  metals,  and  when  beaten  into  very  thin  leaf,  it  appears 
green  by  transmitted  light. 

Gold  is  most  easily  reduced  from  its  combinations.  Most 
metals,  when  placed  in  a  solution  of  a  gold  salt,  precipitate  the 
gold,  and  the  most  feeble  reducing  agents  bring  about  the  same 
result.  On  this  account  a  solution  of  auric  chloride  is  used  for 
toning  photographs.  All  the  compounds  of  gold,  when  ignited  in 
the  air,  are  reduced  to  metallic  gold.  Gold  is  readily  deposited 
upon  other  metals  by  the  process  of  electro-gilding,  the  most 
suitable  solution  being  that  of  the  double  cyanide  of  gold  and 
potassium,  Au(CN)3,KCN. 

Gold  Alloys.  —  Alloys  of  gold  with  copper  and  with  silver 
are  used  for  coinage  and  for  ornamental  purposes,  pure  gold 
being  too  soft  for  these  purposes.  Silver  gives  the  alloy  a  paler 
colour  than  that  of  pure  gold,  while  copper  imparts  to  it  a  reddish 
tinge.  The  alloy  used  for  English  gold  coin  consists  of  gold,  1  1 
parts  ;  copper,  i  part.  The  proportion  of  gold  in  alloys  is  usually 


Gold  Sulphides  569 

expressed  in  parts  per  24  (instead  of  in  percentages),  these  parts 
being  termed  carats.  Thus,  pure  gold  is  said  to  be  24-carat  gold  ; 
i8-carat  gold  contains,  therefore,  18  parts  of  gold  and  6  parts  of 
copper  or  silver.  Most  countries  have  their  own  legal  standards. 
In  England  the  legal  standard  for  gold  coinage  is  22-carats. 

Compounds  of  Gold. — Gold  forms  two  series  of  compounds,  namely,  aurous, 
in  which  the  metal  is  monovalent,  and  auric,  in  which  it  is  trivalent. 

The  composition  of  aurous  compounds  corresponds  to  that  of  the  silver 
compounds.  They  are  very  readily  decomposed.  Thus,  aurous  chloride 
cannot  exist  in  the  presence  of  water,  being  decomposed  into  auric  chloride 
and  metallic  gold.  For  this  reason,  when  aurous  oxide,  Au2O,  is  acted  upon 
by  aqueous  hydrochloride  acid  it  forms  auric,  and  not  aurous  chloride,  thus — 

3Au20  +  6HCl=2AuCl3  +  3H2O  +  4Au. 

With  iodine  geld  forms  only  aurous  iodide,  Aul ;  therefore,  when  auric  oxide 
is  acted  upon  by  hydriodic  acid,  aurous  iodide  and  free  iodine  are  formed, 
thus— 

Au203+6HI=2AuI-j-2I2+3H20. 

Auric  Chloride,  AuCl3,  is  obtained  by  dissolving  gold  in  aqua-regia,  and 
evaporating  the  solution  to  dryness.  When  the  residue  is  dissolved  in  water  the 
concentrated  solution  deposits  reddish  crystals  of  the  composition  AuCl32H2O. 
These  lose  their  water  when  carefully  heated,  leaving  a  brown  mass  of  deliques- 
cent crystals.  Auric  chloride  forms  double  chlorides  with  the  alkaline  chlorides, 
and  with  hydrochloric  acid,  which  may  be  obtained  as  crystalline  compounds. 
Thus,  the  compound  AuCl3HCl,3H2O  is  deposited  from  a  strong  solution  of 
gold  in  aqua-regia.  This  substance  is  sometimes  termed  chloro-auric  acid, 
and  the  double  compounds  with  metallic  chlorides,  such  as  AuCl3,NaCl,2H2O 
and  (AuCl3,KCl)2H2O,  are  known  as  chloro-aurates. 

Auric  Oxide,  Au2O3,  is  obtained  as  a  brown  powder  when  the  hydrated 
oxide,  Au.2O3,3H2O  (or  Au(HO)3),  is  gently  warmed.  At  100°  it  begins  to  de- 
compose, and  at  higher  temperatures  is  completely  converted  into  oxygen  and 
metallic  gold. 

Auric  oxide  is  feebly  basic,  forming  a  few  unstable  salts,  in  which  gold 
replaces  the  hydrogen  in  acids.  It  is  also  a  feeble  acid-forming  oxide,  and 
forms  salts  called  aurates,  such  as  potassium  aurate,  KAuO2,3H2O,  which  may 
be  regarded  as  being  derived  from  an  acid  of  the  composition  HAuO2. 

Auric  oxide  forms  a  compound  with  ammonia,  known  as  fulminating  gold, 
the  exact  composition  of  which  is  not  known.  It  explodes  with  violence  when 
dry  if  struck  or  gently  warmed. 

Gold  Sulphides.— Two  sulphides  of  gold  have  been  obtained,  aurous 
sulphide,  Au^S,  and  auro-auric  sulphide,  Au2S,Au2S3  (or  AuS).  The  latter  is 
formed  when  sulphuretted  hydrogen  is  passed  into  a  cold  solution  of  auric 
chloride — 

8AuCl3 + 9H3S  +  4H20 = 2(  Au2S,  Au2S3)  +  24HC1 + H^O* 


CHAPTER    VI 
ELEMENTS  OF  GROUP  II.  (FAMILY  A.) 


Beryllium,  Be 
Magnesium,  Mg  . 
Calcium,  Ca 

Atomic 
Weights. 
.        9.1 
.      24.32 
.      40.1 

Strontium,  Sr 
Barium,  Ba  . 

Atomic 

Weights. 

87.6 

137-4 


WITH  the  exception  of  the  rare  element  beryllium,  these  metals 
were  first  obtained  (although  not  in  the  pure  state)  by  Davy,  who, 
soon  after  his  discovery  of  the  metals  potassium  and  sodium, 
showed  that  the  so-called  earths  were  not  elementary  bodies  as 
had  been  supposed,  but  were  compounds  of  different  metals  with 
oxygen. 

The  element  beryllium  is  of  later  discovery,  for  although  as 
early  as  1798  it  had  been  shown  by  Vanquelin  that  the  particular 
"  earth  "  in  the  mineral  beryl  was  different  from  any  other  known 
earth,  it  was  not  until  1827  that  the  metal  it  contained  was  iso- 
lated by  Wohler.  In  a  state  approaching  to  purity,  beryllium  was 
first  prepared  by  Humpidge,  1885. 

None  of  the  elements  of  this  family  occurs  in  nature  in  the  un- 
combined  condition  ;  and,  with  the  exception  of  magnesium  and 
calcium,  the  metals  themselves,  in  their  isolated  condition,  are  at 
present  little  more  than  chemical  curiosities.  In  the  case  of 
beryllium  this  is  due  to  the  comparative  rarity  of  its  compounds  ; 
but  with  strontium,  and  barium,  whose  compounds  are  extremely 
abundant,  it  is  owing  partly  to  the  difficulty  of  isolating  the  metals 
in  a  pure  state,  and  also  to  the  fact  that  hitherto  they  have  re- 
ceived no  useful  application.  Beryllium  and  magnesium  are 
white  metals,  which  retain  their  lustre  in  the  air.  Calcium,  stron- 
tium, and  barium  on  exposure  to  air  quickly  become  converted 
into  oxide. 

The  metals  calcium  and  strontium,  as  obtained  by  earlier  experi- 
menters, presented  a  pale  yellow  colour  (it  is  doubtful  whether 

570 


Metals  of  the  Alkaline  Earths  571 

the  metal  barium  was  actually  obtained  by  these  chemists).  But 
the  calcium  which  has  recently  been  obtained  in  considerable 
masses  is  found  to  be  a  silver-white  metal. 

All  these  metals  form  an  oxide  of  the  type  RO.  Beryllium  oxide 
is  insoluble  in  water  ;  magnesium  oxide  is  very  slightly  soluble 
(i  part  in  55,000  or  100,000  parts  of  water),  but  the  solution 
shows  a  feeble  alkaline  reaction.  The  calcium,  strontium,  and 
barium  oxides  show  increasing  solubility,  and  stronger  alkalinity 
and  causticity.  On  this  account  these  elements  are  known  as  the 
metals  of  the  alkaline  earths.  These  three  elements  also  form 
peroxides  of  the  type  RO2. 

All  the  monoxides  are  basic,  and  combine  with  acids  to  form 
salts  of  the  types  RC12,  RSO4,  R(NO3)2. 

The  element  beryllium  (the  typical  element)  stands  apart  from 
the  others  of  this  family  in  many  of  its  chemical  relations.  Thus 
the  oxide  BeO,  unlike  the  corresponding  compounds  of  the  other 
elements,  does  not  combine  with  water  to  form  the  hydroxide. 
The  hydroxide  Be(HO)2  is  soluble  in  sodium  and  potassium 
hydroxide.  In  this  respect  beryllium  exhibits  its  resemblance  to 
zinc.  The  chloride  also  differs  from  the  other  chlorides  in  being 
volatile. 

In  its  permanence  in  air,  its  colour,  its  high  melting-point, 
the  solubility  of  its  sulphate,  and  the  readiness  with  which  its 
hydroxide  is  converted  by  heat  into  the  oxide,  beryllium  ex- 
hibits a  close  similarity  to  magnesium.  In  the  solubility  of 
its  hydroxide  in  potassium  hydroxide,  and  in  its  inability  to 
decompose  water,  beryllium  also  shows  a  marked  resemblance 
to  zinc. 

The  three  elements,  calcium,  strontium,  and  barium,  ex- 
hibit a  closer  resemblance  to  each  other  in  most  of  their 
physical  and  chemical  relations,  than  to  either  magnesium  or 
beryllium. 

They  are  readily  distinguished  by  their  different  spectra. 
Barium  salts,  when  heated  in  a  non-luminous  flame,  impart  to  it 
a  green  colour.  Calcium  and  strontium,  under  the  same  cir- 
cumstances, each  give  a  red  colour  ;  but  the  red  imparted  by 
strontium  compounds  is  more  brilliant,  and  less  orange,  than  that 
of  calcium  salts.  When  the  flames  are  examined  by  the  spectro- 
scope, the  most  characteristic  lines  given  by  barium  are  two  in  the 
bright  green  (Baa  and  Ba/3).  These  are  accompanied  by  a  number 
of  less  brilliant  lines.  The  spectrum  of  strontium  consists  of  four 


572  Inorganic  Chemistry 

specially  prominent  lines,  one  in  the  bright  blue  (Sr<5),  one  in  the 
orange  (Sra),  and  two  in  the  red  (Sr/3  and  Sr-y),  with  others  less 
pronounced  ;  while  that  of  calcium  contains  one  brilliant  green 
line  (Ca/2)  and  one  equally  brilliant  orange  line  (Caa),  besides  a 
large  number  of  less  prominent  lines. 


BERYLLIUM  (Ghtcinnm). 
Symbol  Be.     Atomic  weight  =  9.1. 

Occurrence. — This  element  occurs  principally  in  the  mineral  beryl,  a  double 
silicate  of  the  composition  3BeO,Al,2O3>GSiO2.  The  transparent  varieties  are 
used  as  gems,  the  transparent  green  beryl  being  the  precious  emerald. 

Phenacite  is  beryllium  silicate  Be2SiC>4,  while  chrysoberyl  has  the  compo- 
sition BeO,Al2O3. 

Formation. — The  element  is  obtained  by  heating  sodium  in  the  vapour  of 
beryllium  chloride,  all  air  having  been  previously  replaced  by  hydrogen.  The 
product  is  afterwards  melted  beneath  fused  sodium  chloride,  when  it  is 
obtained  as  a  coherent  solid  metal.  It  may  also  be  obtained  by  the  electro- 
lysis of  the  fused  mixed  chlorides  of  beryllium  and  potassium. 

Properties. — Beryllium  is  a  white  metal  resembling  magnesium.  It  has  a 
specific  gravity  of  2.1,  and  is  moderately  malleable.  It  does  not  readily 
tarnish  in  the  air  at  ordinary  temperatures,  but  when  strongly  heated,  be- 
comes coated  with  a  protecting  film  of  oxide.  The  powdered  metal,  when 
heated,  takes  fire,  and  burns  with  a  bright  light.  It  has  no  action  upon 
water,  even  at  the  boiling  temperature. 

Beryllium  is  easily  dissolved  by  dilute  hydrochloric  acid,  with  evolution  of 
hydrogen.  Cold  dilute  sulphuric  acid  is  without  action,  but  when  heated 
slowly  dissolves  it.  Nitric  acid  slowly  attacks  it  when  concentrated  and  hot. 
It  readily  dissolves  in  potassium  hydroxide,  with  evolution  of  hydrogen. 

Beryllium  Compounds.— The  best  known  are  the  oxide  (berylla],  BeO,  a 
white  infusible  powder,  insoluble  in  water,  soluble  in  acids ;  the  chloride, 
BeCl2,  obtained  by  heating  the  oxide  with  charcoal  in  a  stream  of  chlorine,  a 
white  crystalline  solid,  readily  fused  and  volatilised. 

Beryllium  compounds  do  not  impart  any  colour  to  a  Bunsen  flame.  They 
are  characterised  by  possessing  a  sweet  taste,  hence  the  name  of  glucinum 
originally  given  to  this  element. 


MAGNESIUM. 

Symbol,  Mg,     Atomic  weight =24. 32. 

Occurrence.— Magnesium  is  not  found  in  the  uncombined  state. 
In  combination  it  is  widely  distributed,  and  is  extremely  abundant. 
In  the  mineral  dolomite,  associated  with  calcium  as  carbonate,  it 
occurs  in  mountainous  masses. 

Magnesite,  MgCO3  ;  kieserite^  MgSO4,H2O  ;  carnallite,  MgCl2, 


Magnesium  573 

KC1,6H2O,  are  amongst  the  commoner  naturally  occurring  magne- 
sium compounds.  It  is  also  a  constituent  of  asbestos,  meerschaum, 
serpentine,  talc,  and  a  large  number  of  other  silicates.  As  sulphate 
and  chloride  it  is  met  with  in  sea-water  and  many  saline  springs. 

Modes  of  Formation. — Magnesium  was  obtained  by  Bunsen 
by  the  electrolysis  of  fused  magnesium  chloride  ;  and  later  by 
Matthiessen  by  electrolysing  the  fused  double  chloride  of  mag- 
nesium and  potassium  (carnallite). 

On  a  manufacturing  scale  it  was  later  produced  by  the  reduction 
of  magnesium  chloride  by  means  of  sodium.  A  mixture  of 
anhydrous  magnesium  chloride  (or  fused  mixed  chlorides  of  mag- 
nesium and  sodium,  or  potassium),  powdered  cryolite,  and  sodium 
is  thrown  into  a  red-hot  crucible,  which  is  quickly  closed.  A 
violent  reaction  takes  place,  at  the  conclusion  of  which  the  melted 
mixture  is  stirred  with  an  iron  rod  to  cause  the  globules  of  mag- 
nesium to  run  together. 

The  crude  metal  is  afterwards  purified  by  distillation. 

At  the  present  time  magnesium  is  manufactured  by  a  process 
which  is  practically  that  formerly  employed  by  Matthiessen  on  a 
small  scale,  but  modified  in  detail  to  suit  modern  electrical  re- 
sources. An  iron  crucible  or  melting  pot  is  used,  which  is  made 
the  cathode,  and  the  double  magnesium  potassium  chloride 
(carnallite)  is  maintained  at  a  temperature  about  700° — i.e.  a  dull 
red  heat — by  means  of  gaseous  fuel.  The  anode  consists  of 
a  stout  carbon  rod  which  dips  into  the  molten  material,  and 
is  surrounded  by  a  porcelain  cylinder  which  conveys  away  the 
chlorine. 

Properties. — Magnesium  is  a  silvery-white  metal,  which  does 
not  tarnish  in  dry  air,  but  becomes  coated  with  a  film  of  oxide 
when  exposed  to  air  and  moisture.  At  a  red  heat  it  melts,  and  at 
higher  temperatures  may  be  distilled.  When  heated  in  the  air  it 
takes  fire,  and  burns  with  a  dazzling  white  light,  which  is  extremely 
rich  in  the  chemically  active  rays.  The  flash  of  light,  obtained  by 
projecting  a  small  quantity  of  magnesium  filings  into  a  spirit  flame, 
is  used  for  photographic  purposes.  Magnesium  is  only  moderately 
malleable,  and  is  only  ductile  at  high  temperatures  ;  it  is  readily 
pressed  into  the  form  of  wire  at  a  temperature  slightly  below  its 
melting-point.  Magnesium  only  slightly  decomposes  water  even  at 
the  boiling-point ;  but  when  strongly  heated  in  a  current  of  steam,  the 
metal  takes  fire  (p.  173).  Magnesium  is  rapidly  dissolved  by  dilute 
acids,  with  brisk  evolution  of  hydrogen,  but  solutions  of  caustic 


574  Inorganic  Chemistry 

alkalies  are  unacted  upon  by  it  (compare  Zinc).  When  heated  with 
aqueous  solutions  of  ammonium  salts,  hydrogen  is  evolved,  and  a 
double  salt  of  magnesium  and  ammonium  is  found  in  the  solution. 

Magnesium  combines  directly  with  nitrogen,  when  strongly 
heated  in  that  gas,  forming  magnesium  nitride,  N2Mg3  (p.  232). 

On  account  of  the  brilliant  light  emitted  by  burning  magnesium, 
it  is  employed  for  signalling  purposes,  and  also  in  pyrotechny. 

Magnesium  Oxide  (magnesia},  MgO,  is  found  native  as  the 
mineral  periclase.  It  is  formed  when  magnesium  burns  in  the  air, 
or  when  magnesium  carbonate  is  submitted  to  prolonged  gentle 
calcination,  when  it  is  obtained  as  a  white  bulky  powder,  known  in 
commerce  as  calcined  magnesia  or  magnesia  usta. 

Magnesia  is  extensively  manufactured  from  the  magnesium 
chloride  occurring  in  the  Stassfurt  deposits,  by  first  converting  the 
chloride  into  carbonate  and  subjecting  this  to  calcination.  Mag- 
nesia has  been  obtained  in  the  crystalline  form,  identical  with  that 
of  periclase,  by  heating  the  amorphous  compound  in  a  stream  of 
gaseous  hydrochloric  acid.  It  may  be  fused  in  the  oxyhydrogen 
flame,  and  on  cooling  it  solidifies  to  a  vitreous  mass  which  is  suffi- 
ciently hard  to  cut  glass.  On  account  of  its  extreme  refractoriness, 
magnesia  is  used  for  a  variety  of  metallurgical  purposes,  such  as 
the  manufacture  of  crucibles,  cupels,  &c. 

Magnesium  Hydroxide,  Mg(HO)2,  is  found  in  nature  as  the 
mineral  brucite.  It  is  prepared  by  precipitating  a  magnesium  salt 
by  sodium  or  potassium  hydroxide.  At  a  dull  red  heat  it  loses 
water,  and  is  converted  into  the  oxide,  and  the  magnesia  so 
obtained  has  the  property  of  rehydrating  itself  in  contact  with 
water,  with  evolution  of  heat. 

Magnesium  hydroxide  slowly  absorbs  carbon  dioxide,  forming 
the  carbonate  ;  owing  to  this  fact,  and  to  the  property  it  possesses 
of  rehydration,  magnesia  that  has  been  prepared  by  calcination  at 
a  low  temperature  can  be  employed  as  a  cement.  Thus,  if  calcined 
magnesite  be  made  into  a  paste  with  water,  the  mixture  is  found  to 
harden  in  about  twelve  hours,  and  ultimately  to  acquire  a  hardness 
equal  to  that  of  Portland  cement. 

Magnesium  Chloride,  MgCl2. — This  salt  is  formed  when  mag- 
nesia, or  magnesium  carbonate,  or  the  metal  itself,  is  dissolved  in 
hydrochloric  acid.  From  this  solution  monosymmetric  crystals  of 
the  composition  MgCl2,6H2O  are  deposited.  When  this  salt  is 
heated  it  loses  water,  and  at  the  same  time  is  partially  decomposed 
into  hydrochloric  acid  and  magnesia  ;  in  order,  therefore,  to  pre- 


Magnesium  Sulphate  575 

pare  the  pure  anhydrous  compound,  the  double  magnesium  ammo- 
nium chloride  is  first  formed,  by  adding  ammonium  chloride  to  a 
solution  of  magnesium  chloride.  On  evaporation,  the  double  salt 
separates  out,  MgCl2,NH4Cl,6H2O.  This  salt  allows  itself  to  be 
dehydrated  by  heating,  without  any  decomposition  of  the  magne- 
sium chloride.  When  the  dried  salt  is  more  strongly  heated, 
ammonium  chloride  volatilises  and  leaves  the  anhydrous  magnesium 
chloride  as  a  fused  mass,  which  congeals  to  a  white  crystalline 
solid.  Magnesium  chloride  is  deliquescent,  and  dissolves  in  water 
with  evolution  of  heat.  With  alkaline  chlorides  it  forms  double 
salts,  as  the  ammonium  salt  above  mentioned.  The  potassium 
salt,  MgCl2,KCl,6H2O,  occurs  in  large  quantities  as  the  mineral 
carnalities  and  the  calcium  salt,  2MgCl2,CaCl2,12H2O,  as  tachy- 
drite,  in  the  Stassfurt  deposits.-  When  a  strong  solution  of 
magnesium  chloride  is  made  into  a  thick  paste  with  calcined 
magnesia,  the  mass  quickly  sets  and  hardens,  like  plaster  of  Paris, 
and  is  found  to  contain  an  oxychloride  having  the  composition 
MgCl9,5MgO,  associated  with  varying  quantities  of  water.  The 
white  deposit  which  forms  in  bottles  containing  the  solution  known 
as  magnesia  mixture  consists  of  MgCl2,5MgO,13H2O. 

When  magnesium  oxychloride  is  heated  to  redness  in  a  current 
of  air,  the  magnesium  is  converted  into  oxide,  and  a  mixture  of 
chlorine  and  hydrochloric  acid  is  evolved.  The  reaction  may  be 
represented  as  taking  place  as  follows — 

2MgCl2+H2O  +  O  =  2MgO  +  2HCl  +  Cl2. 

The  Weldon-Pe'chiney  process  for  manufacturing  chlorine  is 
based  upon  this  reaction. 

Magnesium  Sulphate,  MgSO4,7H2O  (Epsom  salts],  is  met  with 
in  many  mineral  springs,  and  in  large  quantities  as  the  mineral 
kieserite,  MgSO4,H2O. 

Magnesium  sulphate  may  be  obtained  by  decomposing  dolomite, 
(CaMg)CO3,  with  sulphuric  acid,  the  nearly  insoluble  calcium 
sulphate  being  readily  removed  from  the  soluble  magnesium  salt 
Magnesium  sulphate  is  now  very  largely  manufactured  from 
kieserite,  which  in  contact  with  water  is  converted  from  the  slightly 
soluble  monohydrated  salt  into  MgSO4,7H2O,  which  is  readily 
soluble,  and  is  purified  by  recrystallisation.  As  usually  obtained, 
crystallised  magnesium  sulphate,  MgSO4,7H2O,  forms  colourless 
rhombic  prisms  ;  but  when  deposited  from  a  cold  supersaturated 


576  Inorganic  Chemistry 

solution,  it  sometimes  forms  prisms  belonging  to  the  monosymmetric 
(monoclinic)  system,  having  the  same  degree  of  hydration.  Above 
50°,  monosymmetric  prisms  of  the  composition  MgSO4,6H2O  are 
deposited. 

When  the  ordinary  salt,  MgSO4,YH2O,  is  placed  over  sulphuric 
acid,  it  loses  two  molecules  of  water  :  when  heated  to  1 50°  it  loses 
six  molecules,  and  at  200°  it  becomes  anhydrous.  At  the  ordinary 
temperature,  100  parts  of  water  dissolve  126  parts  of  crystallised 
magnesium  sulphate  ;  the  solution  has  a  bitter  taste,  and  acts  as  a 
purgative.  With  alkaline  sulphates,  magnesium  sulphate  forms  a 
series  of  double  salts,  having  the  general  formula  MgSO4,R2SO4, 
6H2O.  They  are  ismorphous  with  each  other,  crystallising  in 
monosymmetric  prisms.  The  potassium  salt  occurs  in  the  Stassfurt 
deposits  as  schonite. 

When  anhydrous  magnesium  sulphate  is  dissolved  in  hot  sul- 
phuric acid,  two  acid  sulphates  are  obtained.  One,  having  the  com- 
position MgSO4H2SO4,  is  deposited  from  the  hot  solution  ;  while 
from  the  cold  liquid  the  salt  that  crystallises  has  the  composition 
MgSO4,3H2SO4.  They  are  at  once  decomposed  by  water. 

Magnesium  Carbonate,  MgCO3,  occurs  as  the  mineral  magne- 
stte,  which  is  sometimes  found  as  rhombohedral  crystals,  isomor- 
phous  with  crystals  of  calcite  (CaCO3).  Magnesium  exhibits  a 
great  tendency  to  form  basic  and  hydrated  carbonates  ;  the  normal 
carbonate,  MgCO3,  is  therefore  not  obtained  by  precipitating  a 
magnesium  salt  with  an  alkaline  carbonate  ;  the  white  precipitate 
formed  under  these  circumstances  is  a  basic  carbonate,  whose 
composition  varies  with  the  conditions  of  precipitation.  If,  how- 
ever, this  precipitate  be  suspended  in  water,  and  the  liquid  saturated 
with  carbon  dioxide,  the  compound  dissolves  (more  readily  under 
increased  pressure),  and  when  the  solution  is  heated  to  300°  under 
pressure,  in  such  a  manner  that  the  evolved  carbon  dioxide  can 
escape,  the  normal  anhydrous  carbonate  is  deposited  in  rhombo- 
hedral crystals  isomorphous  with  calcite.  If  the  solution  be 
evaporated  to  dryness,  the  normal  carbonate  is  deposited  in 
rhombic  crystals  isomorphous  with  arragonite  (CaCO3).  Magne- 
sium and  calcium  carbonates  are  therefore  isodimorphous. 

Basic  Carbonates. — The  mineral  hydromagnesite  is  a  basic 
carbonate  of  the  composition  3MgCO3,Mg(HO)2,3H2O.  A  number 
of  basic  carbonates  are  formed  by  the  precipitation  of  a  magnesium 
salt  with  sodium  carbonate.  Thus,  under  ordinary  conditions  a 
white  bulky  precipitate  is  obtained,  known  in  pharmacy  as  magnesia 


Calcium  577 

alba  levis.  Its  composition,  although  liable  to  vary  through  the 
presence  of  other  basic  carbonates,  is  in  the  main  the  same  as  that 
of  hydromagnesite. 

If  the  precipitation  be  made  with  boiling  solutions,  and  the  pre- 
cipitate so  obtained  be  dried  at  100°,  a  denser  carbonate  is  ob- 
tained, termed  magnesia  alba ponderosa,  4MgCO3>Mg(HO)2,4H2O. 

When  an  excess  of  sodium  carbonate  is  employed,  and  the 
mixture  is  subjected  to  prolonged  boiling,  a  carbonate  is  obtained 
having  the  composition  2MgCO3,Mg(HO)2,2H2O. 


CALCIUM. 

Symbol,  Ca.     Atomic  weight  =  40.1. 

Occurrence.  —  Calcium  is  only  met  with  in  nature  in  combina- 
tion. It  occurs  in  enormous  quantities  as  the  carbonate  in  a  great 
variety  of  different  minerals,  such  as  marble,  limestone,  calcspar, 
and  also  as  coral;  and  with  carbonate  of  magnesium  as  dolomite, 
or  magnesian  limestone.  In  the  form  •  of  sulphate,  calcium 
occurs  as  gypsum  and  selenite,  CaSO4,2H2O,  and  as  anhydrite, 
CaSO4.  The  fluoride  CaF2  occurs  as  fluorspar,  and  the  various 
silicious  rocks  contain  compound  silicates  of  calcium  and  other 
metals.  The  carbonate  and  sulphate  are  present  in  most  spring 
and  river  waters.  Calcium  compounds  are  also  present  in  all 
vegetable  and  animal  organisms.  Thus,  bones  consist  largely  of 
calcium  phosphate. 

Modes  of  Formation.  —  Although  calcium  compounds  are  so 
extremely  abundant,  the  metal  itself,  until  quite  recently,  was 
scarcely  more  than  a  chemical  curiosity.  The  element  was  first 
isolated  in  an  impure  state  by  Davy  (1808). 

More  recently  Moissan  obtained  the  metal  in  the  form  of  crystals 
by  heating  together  sodium  and  calcium  iodide  — 


The  calcium  dissolves  in  the  excess  of  sodium,  and  on  cooling  it  cry- 
stallises out.  The  sodium  is  removed  by  solution  in  absolute  alcohol. 
At  the  present  time  calcium  is  obtained  commercially  by  the 
electrolysis  of  the  fused  chloride,  the  success  of  the  process  de- 
pending upon  the  device  adopted  for  removing  the  metal,  as  it  is 
reduced,  from  the  action  of  the  fused  electrolyte.  The  cathode 
employed  is  a  rod  of  iron  which  is  brought  just  to  the  surface  of 
the  melted  chloride.  As  soon  as  a  small  quantity  of  the  metal 
calcium  collects  beneath  the  end  of  the  cathode,  the  latter  is  very 

2o 


c^8  Inorganic  Chemistry 

slowly  raised  by  a  suitable  mechanical  arrangement,  so  that  the 
calcium  may  solidify  upon  the  end  of  the  iron  rod  without  any 
interruption  of  the  electrolysis.  As  this  process  of  continuously 
raising  the  cathode  proceeds,  a  rugged  rod  or  bar  of  calcium 
weighing  several  pounds  may  be  gradually  built  up. 

Properties.— Calcium  is  a  silver-white  metal  having  a  specific 
gravity  1.85,  and  melting  about  760°  C.  It  is  moderately  soft  and 
malleable.  The  metal  is  readily  oxidised  by  moist  air,  and 
decomposes  water  at  the  ordinary  temperature.  When  heated 
in  air  it  takes  fire  and  burns.  When  heated  in  hydrogen  it 
forms  calcium  hydride,  CaH2. 

Oxides  of  Calcium. — Two  oxides  are  known,  namely,  calcium 
monoxide,  CaO,  and  calcium  dioxide,  CaO2. 

Calcium  Oxide  (time,  quicklime),  CaO,  is  obtained  by  heating 
calcium  carbonate  to  redness — 

CaCO3  =  CO2  +  CaO. 

On  a  large  scale  lime  is  manufactured  by  burning  limestone  or 
chalk  in  kilns  with  coal.  If  much  clay  be  present  with  the  lime- 
stone, care  is  required  to  prevent  the  mass  from  fusing  when  it  is 
said  to  be  dead  burnt.  Lime  is  a  white  amorphous  substance, 
which  is  infusible  by  the  oxyhydrogen  flame,  but  which,  when  so 
heated,  emits  a  bright  light,  known  as  the  oxyhydrogen  limelight. 
It  absorbs  moisture  and  carbon  dioxide  from  the  air.  On  account 
of  its  power  of  absorbing  moisture,  lime  is  frequently  employed  as 
a  dehydrating  agent.  Thus,  gases  which  cannot  be  dried  by  means 
of  sulphuric  acid  (e.g.  ammonia)  may  be  deprived  of  moisture  by 
being  passed  over  calcium  oxide.  It  is  also  used  for  withdrawing 
water  from  alcohol  in  the  preparation  of  absolute  alcohol.  When 
a  small  quantity  of.  water  is  poured  upon  lime  the  mass  rapidly 
becomes  hot,  and  volumes  of  steam  are  given  off,  the  lime  at  the 
same  time  swelling  up  and  crumbling  to  a  soft,  dry  powder.  This 
process  is  known  as  the  slaking  of  lime,  and  the  product  is  termed 
slaked  lime^  in  contradistinction  to  quick  lime.  The  lime  enters 
into  chemical  union  with  water,  forming  calcium  hydroxide,  thus — 

CaO  +  H2O  =  Ca(HO)2. 

Calcium  Hydroxide,  Ca(HO)2,  is  a  white  amorphous  powder, 
sparingly  soluble  in  water,  and,  unlike  the  majority  of  solids,  it  is 
less  soluble  in  hot  than  in  cold  water.  100  parts  of  water  at  the 


Calcium  Chloride  5/9 

ordinary  temperature  dissolve  0.14  part  of  calcium  hydroxide, 
while  at  100°  the  same  volume  of  water  dissolves  about  half  that 
amount.  This  solution,  known  as  lime-water,  has  an  alkaline 
reaction,  and  absorbs  carbon  dioxide,  with  the  precipitation  of 
calcium  carbonate. 

Milk  of  Lime  is  the  name  given  to  a  mixture  of  lime  with  less 
water  than  will  dissolve  it,  whereby  an  emulsion  of  lime  is  obtained. 
When  a  thick  paste  of  lime  and  water  is  exposed  to  the  atmos- 
phere, in  a  few  days  it  sets,  and  continues  gradually  to  harden. 
On  this  account  lime  is  used  for  mortars  and  cements.  Mortar 
consists  of  a  mixture  of  lime  and  sand  with  water.  The  sand 
serves  the  double  purpose  of  preventing  shrinkage  on  drying,  and 
also  of  rendering  the  mass  more  permeable  to  atmospheric  carbon 
dioxide.  The  setting  of  mortar  is  due  to  the  combined  action  of 
evaporation  and  absorption  of  carbon  dioxide. 

Calcium  Dioxide,  CaO2,  is  obtained  by  adding  lime-water  to 
hydrogen  peroxide,  or  to  sodium  peroxide  acidulated  with  dilute 
nitric  acid  ;  sparingly  soluble  crystals  of  CaO2,8H2O  separate  out, 
which  at  130°  lose  their  water.  When  more  strongly  heated  the 
monoxide  is  formed  with  evolution  of  oxygen. 

Calcium  Chloride,  CaCl2,  occurs  in  sea  and  river  waters, 
and  is  present  in  tacky drite  of  the  Stassfurt  deposits.  It 
is  obtained  in  large  quantities  as  a  bye-product  in  many  manu- 
facturing processes,  such  as  that  of  potassium  chlorate,  ammonia 
from  ammonium  chloride,  &c.  It  may  be  obtained  by  the  action 
of  hydrochloric  acid  upon  calcium  carbonate,  and  is  deposited  on 
concentration,  in  large  colourless,  deliquescent,  hexagonal  prisms, 
CaCl2,6H2O,  which  melt  at  29°  in  their  water  of  crystallisation. 
When  heated  below  200°  the  crystals  part  with  four  molecules  of 
water,  and  above  200°  become  anhydrous.  As  thus  obtained  the 
anhydrous  salt  is  a  porous  mass,  which  is  extremely  hygroscopic, 
and  on  this  account  is  used  as  a  desiccating  agent,  both  for  gases 
and  liquids.  At  a  red  heat  it  fuses,  and  on  cooling  solidifies 
to  a  crystalline,  deliquescent  mass.  Calcium  chloride  combines 
with  ammonia,  forming  the  compound  CaCl2,8NH3.  Calcium 
chloride,  therefore,  cannot  be  employed  for  drying  gaseous 
ammonia. 

Crystallised  calcium  chloride  is  extremely  soluble  in  water ; 
100  parts  of  water  at  16°  dissolve  400  parts  of  the  salt,  the  solu- 
tion being  attended  with  considerable  absorption  of  heat.  When 
mixed  with  powdered  ice  or  snow  liquefaction  of  both  the  solids 


580  Inorganic  Chemistry 

rapidly  takes  place,  and  the  consequent  absorption  of  heat  lowers 
the  temperature  of  the  mixture  to  -  40°. 

Bleaehing-Powder  (chloride  of  time),  Ca(OCl)Cl.  —  This  im- 
portant compound  is  manufactured  on  a  large  scale  by  the  action 
of  chlorine  upon  slaked  lime.  The  hydrated  lime  is  spread  upon 
the  floor  of  the  bleaching-powder  chambers  to  a  depth  of  three  or 
four  inches,  and  raked  into  ridges  or  furrows  with  a  special  wooden 
rake.  Chlorine  is  then  led  into  the  chambers,  which  are  provided 
with  glass  windows  to  enable  the  operator  to  examine  the  colour 
of  the  atmosphere  within.  At  first  the  absorption  of  the  chlorine  is 
rapid,  but  as  the  reaction  proceeds  it  becomes  slower,  and  the  lime 
is  from  time  to  time  raked  over  to  expose  a  fresh  surface.  The 
lime  is  left  in  contact  with  the  gas  for  twelve  to  twenty-four  hours. 
The  excess  of  chlorine  is  absorbed  by  projecting  into  the  chamber 
a  shower  of  fine  lime  dust  by  means  of  a  mechanical  fan-distributor. 
This,  in  settling,  rapidly  absorbs  all  the  chlorine,  and  the  chambers 
can  then  be  opened  without  any  unpleasant  smell  of  chlorine  being 
perceptible. 

The  reaction  which  takes  place  is  expressed  by  the  equation  — 

Ca(HO)2  +  Cl2==Ca(OCl)Ci  +  H2O. 

It  was  formerly  believed  that  bleaching-powder  was  a  mechani- 
cal mixture  of  calcium  chloride,  CaCl2,  and  calcium  hypochlorite, 
Ca(OCl)2,  but  it  has  been  conclusively  shown  that  the  substance 
does  not  contain  any  free  calcium  chloride.  It  may,  however,  be 
regarded  as  a  compound  consisting  of  equivalent  proportions  of 
these  two  salts,  and  its  composition  may  be  expressed  by  the  for- 
mula Ca(OCl)2,CaCl2,  which  corresponds  to  2Ca(OCl)Cl. 

The  relation  in  which  bleaching-powder  stands  to  calcium  chlo- 
ride on  the  one  hand  and  calcium  hypochlorite  on  the  other  will 
be  seen  by  the  following  formulae  — 


Calcium  Chloride.  Calcium  Hypochlorite.         c 

Cl—  Ca—  Cl  CIO—  Ca—  OC1  Cl—  Ca—  OC1. 

In  practice  the  absorption  of  chlorine  by  the  lime  is  never  as 
complete  as  is  represented  by  the  above  equation,  and  the  com- 
mercial value  of  the  product  depends  upon  the  amount  of  available 
chlorine  it  contains,  i.e.  chlorine  which  is  evolved  on  treating  the 
compound  with  hydrochloric  or  sulphuric  acid.  This  ranges  from 
30  to  38  per  cent. 


Plaster  of  Paris  581 

When  treated  with  water,  bleaching-powder  is  converted  into 
calcium  chloride  and  hypochlorite,  thus  — 

2Ca(OCl)Cl  =  CaCl2  +  Ca(OCl)2. 

Bleaching-powder  decomposes  slowly  even  in  stoppered  bottles, 
and  more  rapidly  on  exposure  to  atmospheric  moisture  and  carbon 
dioxide. 

When  acted  upon  by  acids  chlorine  is  evolved,  thus  — 


Ca(OCl)Cl  +  2HCl    =  CaCl2  +H2O  +  C12 
Ca(OCl)Cl  +  H2SO4  =  CaSO4  +  H2O  +  C12. 

When  a  solution  of  bleaching-powder  is  treated  with  very  dilute 
acids,  hypochlorous,  acid  is  first  liberated,  which  in  contact  with 
hydrochloric  acid  yields  chlorine  — 


(1)  Ca(OCl)2.Aq  +  2HCl.Aq  =  CaCl2  +  2HClO.Aq. 

(2)  HC1O  +  HC1  =  H20  +  C12. 

In  the  process  of  bleaching,  the  material  is  first  steeped  in  a 
dilute  solution  of  bleaching-powder  and  then  in  dilute  acid.  The 
hypochlorous  acid  first  formed  is  decomposed  in  the  presence  of 
excess  of  hydrochloric  acid,  generating  chlorine  within  the  fibres 
of  the  wet  cloth. 

Calcium  Sulphate,  CaSO4,  occurs  as  the  mineral  anhydrite, 
and  in  the  hydrated  condition  as  gypsum,  CaSO4,2H2O,  of  which 
satinspar  (or  fibrous  gypsum},  alabaster,  and  selenite  are  different 
varieties.  It  is  obtained  in  the  hydrated  condition  by  precipita- 
tion from  a  solution  of  calcium  chloride,  on  the  addition  of  sul- 
phuric acid  or  a  soluble  sulphate.  When  dried  at  110°  to  120°  it 
loses  a  portion  of  its  water,  leaving  the  hydrate,  (CaSO4)2,H2O  ;  at 
200°  it  becomes  anhydrous.  Calcium  sulphate,  in  the  hydrated 
condition,  is  slightly  soluble  in  water,  the  solubility  reaching  a 
maximum  at  35°,  when  I  part  of  the  compound  requires  432  parts 
of  water  for  its  solution  ;  above  this  temperature  the  solubility 
again  diminishes.  Its  solubility  is  increased  by  the  presence  of 
alkaline  chlorides  and  free  hydrochloric  acid. 

When  boiled  in  strong  sulphuric  acid  calcium  sulphate  partially 
dissolves,  and  on  cooling  an  acid  sulphate  crystallises  out,  having 
the  composition  CaSO4,H2SO4. 

Plaster  of  Paris  is  calcium  sulphate  which  has  been  partially 
deprived  of  its  water  of  hydration  by  heat,  and  converted  into  the 


582  Inorganic  Chemistry 

hydrate,  (CaSO4)2,H2O.  It  is  manufactured  by  burning  gypsum  in 
a  kiln  or  oven  in  such  a  way  that  the  carbonaceous  fuel  does 
not  come  in  contact  with  the  sulphate,  which  would  result  in  its 
reduction  to  sulphide  ;  the  temperature  is  not  allowed  to  exceed 
about  130°.  If  heated  more  strongly  (above  200°)  the  sulphate 
becomes  anhydrous,  and  is  said  to  be  dead  burnt;  in  this  con- 
dition its  property  of  setting  when  mixed  with  water  is  greatly 
impaired.  When  plaster  of  Paris  is  made  into  a  paste  with  water 
it  rapidly  sets  to  a  hard  mass  ;  this  setting  is  due  to  its'  rehydra- 
tion,  whereby  gypsum  is  reformed,  thus — 

(CaSO4)2,H20  +  3H2O  -  2CaSO4,2H2O  •» 

Calcium  Carbonate,  CaCO3. — This  compound  is  extensively 
met  with  in  nature,  as  limestone,  chalk,  marble,  and  innumerable 
varieties  of  calcspar.  It  is  formed  when  lime  is  exposed  to  atmos- 
pheric carbon  dioxide.  It  is  obtained  when  an  alkaline  carbonate 
is  added  to  a  soluble  calcium  salt. 

Calcium  carbonate  is  dimorphous  ;  it  occurs  as  arragonite  in 
crystals  belonging  to  the  orthorhombic  system,  and  as  calcspar  in 
crystals  belonging  to  the  hexagonal  system.  Both  these  crystal- 
line varieties  can  be  artificially  obtained  ;  when  deposited  from 
solutions  at  the  ordinary  temperature  the  crystals  are  identical 
with  calcite  ;  but  when  crystallised  from  hot  solutions,  they  form 
rhombic  crystals  corresponding  to  arragonite. 

Calcium  carbonate  is  nearly  insoluble  in  water  ;  1000  grammes 
of  water  dissolve  .0018  gramme  of  the  compound.  It  is  more 
soluble  in  water  charged  with  carbon  dioxide,  forming  the  acid 
carbonate  of  lime,  CaCO3,H2CO3,  or  H2Ca(CO3)2. 

looo  grammes  of  water  saturated  with  carbon  dioxide  will  dis- 
solve, at  o°,  0.7  gramme  of  calcium  carbonate.  By  increasing  the 
pressure  (thereby  increasing  the  amount  of  dissolved  gas)  as  much 
as  3  grammes  of  calcium  carbonate  may  be  dissolved.  When  this 
solution  is  boiled  the  acid  carbonate  is  decomposed  (page  221). 

Calcium  Phosphate  (tricalcium  orthophosphate),  Ca3(PO4)2,  is 
the  most  important  of  the  phosphates  of  calcium.  It  is  found  as 
the  mineral  osteolite,  Ca3(PO4)2,2H2O,  and  also  as  sombrerite, 
estramadurite,  and  coprolites.  Apatite  consists  of  phosphate  and 
fluoride,  3Ca3(PO4)2,CaF2 ;  and  the  mineral  constituents  of  bones 
consist  chiefly  of  calcium  phosphate. 

It  is  obtained  in  a  pure  state  by  the  addition  of  ordinary  sodium 


Calcium  Sulphide  583 

phosphate  to  a  solution  of  calcium  chloride  in  the  presence  of 
ammonia.  The  precipitate  is  decomposed  on  boiling  into  an 
insoluble  basic  salt  and  a  soluble  acid  salt.  Although  nearly 
insoluble  in  pure  water,  calcium  phosphate  dissolves  in  water  con- 
taining salts  in  solution,  such  as  sodium  chloride  or  nitrate,  or 
even  dissolved  carbon  dioxide.  On  this  fact  depends  the  readi- 
ness with  which  this  substance  is  absorbed  by  the  roots  of  plants. 
Calcium  phosphate  is  readily  soluble  in  both  nitric  and  hydro- 
chloric acids.  It  is  decomposed  by  sulphuric  acid,  with  the  forma- 
tion of  monocalcium  orthophosphate  and  calcium  sulphate,  thus  — 

Ca3(PO4)2  +  2H  S2O4  =  2CaSO4  +  H4Ca(PO4)2. 

This  mixture  of  calcium  sulphate  and  monocalcium  phosphate 
is  known  as  superphosphate  of  lime,  and  is  largely  used  as  an  arti- 
ficial manure. 

With  a  larger  quantity  of  sulphuric  acid  the  phosphate  is  con- 
verted into  tribasic  phosphoric  acid.  (See  Phosphorus,  page  453.) 

Calcium  Carbide,  CaC2.—  This  compound  is  produced  when 
lime  or  chalk  is  heated  with  carbon  in  the  electric  furnace.  It  is 
also  obtained  as  a  second  product  in  the  manufacture  of  phosphorus 
when  calcium  phosphate  is  heated  with  carbon  (see  Phosphorus). 
Calcium  carbide  is  manufactured  on  an  extensive  scale  for  use  in 
the  preparation  of  acetylene  (page  318). 

Calcium  Sulphide,  CaS,  is  formed  when  sulphuretted  hydrogen 
is  passed  over  heated  lime  — 


Ca(HO)2+H2S 
Or  by  heating  calcium  sulphate  with  carbon  — 
CaSO4 


Calcium  sulphide  is  decomposed  on  boiling  with  water,  forming 
calcium  hydroxide  and  hydrosulphide,  thus  — 

2CaS  +  2H2O  =  Ca(HO)2  +  Ca(HS)2. 

Calcium  sulphide  (in  common  with  barium  and  strontium  sul- 
phides), as  usually  obtained,  possesses  the  property  of  emitting  a 
feeble  light  (or  phosphorescence)  in  the  dark,  after  being  previously 
exposed  to  a  bright  light.  The  light  emitted  gradually  diminishes 
in  intensity,  but  on  re-exposing  the  compound  to  the  light  its 


584  Inorganic  Chemistry 

luminosity  is  again  restored.  This  property  has  been  long  known, 
and  calcium  sulphide  was  formerly  termed  Cantoris  phosphorus. 
The  material  formerly  known  as  Bononian  (or  Bologniari)  phos- 
phorus is  the  corresponding  barium  compound. 

These  various  sulphides  are  now  manufactured  for  the  preparation  of 
so-called  luminous  paint.  The  phosphorescence  of  these  compounds  appears 
to  be  due  to  the  presence  of  small  quantities  of  foreign  substances ;  thus,  not 
only  is  the  particular  colour  of  the  light  emitted  changed  by  the  intentional 
introduction  of  minute  traces  of  bismuth,  cadmium,  manganese,  zinc,  and 
many  other  metals,  but  it  has  been  shown,  in  the  case  of  calcium  sulphide, 
that  the  perfectly  pure  substance  does  not  exhibit  phosphorescence. 


STRONTIUM. 

Formula,  Sr.     Atomic  weight =87. 6. 

Occurrence. — The  chief  natural  compounds  of  this  element  are 
strontianite,  SrCO3,  and  celestine,  SrSO4. 

Modes  of  Formation. — The  metal  was  first  obtained  in  small 
quantity  by  Davy,  by  the  electrolysis  of  the  hydroxide,  or  chloride, 
moistened  with  water. 

It  is  more  advantageously  obtained  by  electrolysing  the  fused 
chloride.  Pure  strontium  has  recently  been  prepared*  by 
strongly  heating  strontium  hydride,  SrH2,  in  vacuo. 

Properties. — Strontium  is  a  silver- white  metal  which  melts 
about  800°.  It  is  readily  oxidised  by  air,  and  decomposes  water 
at  ordinary  temperatures.  When  heated  in  hydrogen  it  forms 
strontium  hydride,  SrH2.  At  —60°  it  combines  with  dry  ammonia, 
forming  red-brown  crystals  of  a  compound  known  as  "  strontium- 
ammonium,"  believed  to  have  the  composition  Sr,6NH3. 

Oxides  Of  Strontium. — Two  oxides,  corresponding  to  those  of 
calcium,  are  known,  namely,  strontium  monoxide,  SrO,  and  dioxide, 
SrO2. 

Strontium  Monoxide  (strontia},  SrO,  is  obtained  by  heating  the 
nitrate  or  carbonate.  It  is  prepared  on  a  large  scale  by  decompos- 
ing strontium  carbonate  by  superheated  steam  ;  carbon  dioxide  is 
evolved,  and  strontium  hydroxide  remains,  which  on  ignition  forms 
the  monoxide.  Strontia  strongly  resembles  lime.  When  treated 
with  water  it  slakes  with  evolution  of  heat,  forming  strontium 
hydroxide,  Sr(HO)2.  The  hydroxide  is  more  soluble  in  water  than 
the  lime  compound,  and  the  solution  on  cooling  deposits  tetragonal 
crystals,  Sr(HO)2,8H2O.  The  solution  is  strongly  alkaline. 

*  Guntz  and  Roederer,  Compt.  Rend. ,  1906. 


Strontium  Nitrate  585 

Strontium  hydroxide  combines  with  sugar,  forming  a  saccharate 
of  strontia,  which  is  readily  decomposed  by  carbon  dioxide.  On 
this  account  it  is  prepared  on  a  large  scale  for  use  in  the  manu- 
facture of  beet-sugar.  One  process  by  which  it  is  obtained  on  a 
commercial  scale  consists  in  first  forming  strontium  sulphide,  by 
reducing  the  natural  sulphate  with  carbon,  and  treating  the  solution 
of  the  sulphide  with  sodium  hydroxide,  thus — . 

SrS  +  NaHO  +  H2O  =  Sr(HO)2  +  NaHS. 

Strontium  Dioxide,  SrO2. — When  hydrogen  peroxide  is  added 
to  a  solution  of  strontium  hydroxide,  a  hydrate  of  the  peroxide 
separates  out  in  the  form  of  pearly  crystals,  SrO2,8H2O.  On  gently 
heating  this  compound,  it  is  converted  into  the  anhydrous  peroxide. 
On  heating  to  redness  it  evolves  oxygen,  and  is  converted  into  the 
monoxide. 

Strontium  Chloride,  SrCl2,  is  obtained  from  strontianite  by  the 
action  of  hydrochloric  acid.  The  salt  deposits  from  the  solution  in 
deliquescent  hexagonal  prisms,  SrCl2,6H2O,  isomorphous  with  the 
corresponding  calcium  compound. 

Strontium  Sulphate,  SrSO4. — The  native  compound  celestine 
occurs  in  amorphous  fibrous  masses,  and  also  in  rhombic  crystals. 
The  name  of  the  mineral  is  derived  from  the  fact  that  it  usually 
has  a  light  blue  coloutf  It  is  produced  by  precipitation  from  a 
strontium  salt  by  sulphuric  acid.  It  is  only  slightly  soluble  in  cold 
water,  and  still  less  in  hot.  When  boiled  with  solutions  of  alkaline 
carbonates,  strontium  sulphate  is  completely  converted  into  stron- 
tium carbonate — 

SrS04  +  Na2CO3  =  SrCO3  +  Na2SO4. 

In  this  respect  strontium  sulphate  differs  from  barium  sulphate, 
which  under  these  conditions  remains  unchanged.  On  treatment 
with  strong  sulphuric  acid,  strontium  sulphate  forms  SrSO4,H2SO4, 
which,  like  the  corresponding  calcium  compound,  is  converted  by 
water  into  sulphuric  acid  and  the  normal  sulphate. 

Strontium  Nitrate,  Sr(NO3)2,  is  obtained  by  dissolving  the 
natural  carbonate  in  dilute  nitric  acid.  On  concentration,  the 
anhydrous  salt  separates  out  in  octahedrons.  From  dilute  solu- 
tion, on  cooling,  it  forms  monosymmetric  prisms,  Sr(NO3)2, 
4H2O,  which  effloresce  on  exposure  to  the  air.  When  heated  with 
carbon,  or  other  readily  combustible  substances,  the  mixture  in- 


586  Inorganic  Chemistry 

flames  and  burns  with  the  red  colour  characteristic  of  strontium 
compounds  ;  strontium  nitrate  is  therefore  largely  used  in  pyro- 
techny  for  the  production  of  red  fire.  This  property  is  most 
readily  illustrated  by  mixing  dry  powdered  strontium  nitrate  with 
ammonium  picrate,  and  igniting  the  mixture,  which  burns  with  a 
brilliant  red  light. 


BARIUM. 

Symbol,  Ba.     Atomic  weight =137. 4. 

Occurrence. — The  most  abundant  natural  compounds  of  barium 
are  heavy  spar,  BaSO4,  and  witherite^  BaCO3.  It  occurs  also, 
associated  with  calcium,  in  the  mineral  barytocalcite^  BaCO3,CaCO3. 

Modes  Of  Formation.— The  element  barium  is  more  difficult  to  isolate  than 
either  strontium  or  calcium,  and  it  is  doubtful  whether  pure  barium  has  ever 
been  obtained.  Davy  electrolysed  various  barium  salts,  made  into  a  thick 
paste  with  water,  using  mercury  as  the  negative  electrode :  in  this  way  an 
amalgam  of  barium  was  formed,  from  which,  on  distilling  away  the  mercury, 
a  dark  porous  mass  was  obtained.  Amalgams  of  barium  and  mercury  have 
been  prepared  in  other  ways,  but  it  has  .been  shown  that  the  product  obtained 
after  distilling  the  mercury  from  these  is  not  pure  barium,  but  is  a  solid  alloy 
or  compound  of  barium  with  mercury. 

By  the  electrolysis  of  the  fused  chloride,  Matthiessen  obtained  small  globules 
of  metal,  which  on  exposure  to  the  air  rapidly  oxidised.  More  recent  experi- 
menters fail  to  obtain  the  metal  by  this  process  (Limb.,  Compt,  Rend.,  112). 

Oxides  of  Barium. — Two  oxides  are  known,  namely,  barium 
monoxide,  BaO,  and  dioxide,  BaO2. 

Barium  Monoxide  (baryta\  BaO,  is  usually  prepared  by  heat- 
ing the  nitrate.  The  mass  fuses  and  evolves  oxygen  and  oxides  of 
nitrogen,  leaving  a  greyish  white  friable  residue  of  the  oxide.  It 
may  also  be  obtained  by  heating  the  carbonate  ;  but  as  the  tem- 
perature necessary  to  expel  the  carbon  dioxide  is  very  high,  it  is 
usual  to  mix  the  carbonate  with  lampblack,  tar,  or  other  sub- 
stances which  on  heating  will  yield  carbon,  when  the  conversion 
takes  place  more  readily,  carbon  monoxide  being  evolved,  thus — 

BaCO3  +  C  =  BaO  +  SCO. 

Small  quantities  may  readily  be  obtained  by  heating  barium 
iodate  in  a  porcelain  crucible,  when  the  iodate  is  decomposed  as 
follows — 


Barium  Dioxide  587 

Barium  oxide  is  a  strongly  caustic  and  alkaline  compound  ;  in 
contact  with  water  it  slakes  with  evolution  of  so  much  heat  that 
the  mass  may  become  visibly  red  hot  if  too  much  water  be  not 
added. 

When  heated  to  a  dull  red  heat  in  oxygen,  or  air,  it  takes  up  an 
additional  atom  of  oxygen  and  forms  the  dioxide  (see  p.  184). 

Barium  Hydroxide,  Ba(HO)2,  is  obtained  when  the  monoxide 
is  slaked  with  water.  It  is  manufactured  by  first  heating  the 
powdered  native  sulphate  with  coal,  when  a  crude  barium  sulphide 
is  formed.  This  is  then  heated  in  a  stream  of  moist  carbon 
dioxide,  whereby  it  is  converted  into  the  carbonate,  and  super- 
heated steam  is  then  passed  over  the  heated  carbonate  — 


BaS  +  H2O  +  CO2  =  BaCO3+  H2S. 
BaC03  +  H2O  =  Ba(HO)2  +  CO2. 

Barium  hydroxide  is  soluble  in  water  :  the  solution,  known  as 
baryta-water^  absorbs  carbon  dioxide,  with  the  precipitation  of 
barium  carbonate. 

The  aqueous  solution  deposits  .  crystals  of  hydrated  barium 
hydroxide,  Ba(HO)2,8H2O,  in  the  form  of  colourless  tetragonal 
prisms,  which  on  exposure  to  the  air  lose  seven  molecules  of  water. 

Barium  hydroxide,  when  heated  in  a  current  of  air,  yields  barium 
dioxide. 

Barium  hydroxide  was  formerly  employed  in  sugar-refining,  but 
owing  to  its  poisonous  nature  it  has  been  superseded  by  strontium 
hydroxide  (q.v.\ 

Barium  Dioxide  (barium  peroxide),  BaO2.  —  This  oxide  is 
obtained  by  heating  the  monoxide  to  a  low  red  heat  in  a  stream  of 
oxygen,  or  of  air  which  has  been  deprived  of  atmospheric  carbon 
dioxide. 

The  pure  compound  may  be  obtained  by  adding  an  excess  of 
baryta-water  to  hydrogen  peroxide,  when  hydrated  barium  per- 
oxide separates  out  in  crystalline  scales  — 

Ba(HO)2+H202  +  6H20  =  Ba02,8H,0. 

On  drying  in  vacuo  at  130°  this  compound  loses  water  and  is 
converted  into  the  anhydrous  peroxide. 

The  commercial  peroxide  may  be  purified  by  treatment  with 
dilute  hydrochloric  acid,  whereby  barium  chloride  and  hydrogen 
peroxide  are  formed.  After  the  removal  of  insoluble  impurities  by 


588  Inorganic  Chemistry 

filtration,  baryta-water  is  cautiously  added,  which  causes  the  pre- 
cipitation of  ferric  oxide  and  silica.  The  liquid  is  then  filtered, 
and  to  the  clear  liquid,  consisting  of  a  solution  of  barium  chloride 
and  hydrogen  peroxide,  an  excess  of  strong  baryta-  water  is  added, 
when  the  hydrated  barium  peroxide  is  precipitated,  as  already 
explained. 

Barium  peroxide  is  a  grey  powder,  which  on  being  heated  to  a 
bright  red  heat  gives  up  oxygen  and  forms  the  monoxide  (p.  184). 

Dilute  acids  decompose  barium  peroxide,  with  formation  of 
hydrogen  peroxide  and  a  barium  salt.  Concentrated  sulphuric 
acid  forms  barium  sulphate  and  ozonised  oxygen.  When  gently 
warmed  in  a  stream  of  sulphur  dioxide,  the  mass  becomes  incan- 
descent and  forms  barium  sulphate  — 

BaO2  +  SO2  =  BaSO4. 

Barium  Chloride,  BaCl2,  may  be  obtained  by  dissolving  the 
natural  carbonate  in  hydrochloric  acid.  It  may  be  obtained  from 
the  natural  sulphate,  either  by  first  converting  it  into  the  sulphide, 
and  decomposing  that  with  hydrochloric  acid,  or  by  roasting  the 
mineral  with  powdered  coal,  limestone,  and  calcium  chloride,  when 
the  following  reactions  take  place  — 


The  barium  chloride  is  dissolved  in  water,  and  an  insoluble  oxy- 
sulphide  of  calcium  remains. 

Barium  chloride  forms  colourless  rhombic  tables,  BaCl2,2H2O, 
which  at  15.6°  are  soluble  to  the  extent  of  43.5  parts  in  100  parts 
of  water.  The  salt  is  nearly  insoluble  in  hydrochloric  acid,  and 
may  therefore  be  precipitated  from  an  aqueous  solution  by  the 
addition  of  this  acid. 

Barium  chloride,  in  common  with  all  the  soluble  salts  of  this 
element,  is  highly  poisonous. 

Barium  Sulphate,  BaSO4,  is  the  most  abundant  naturally 
occurring  barium  compound.  It  is  frequently  met  with  as  large 
rhombic  crystals.  The  specific  gravity  of  the  mineral  is  4.3  to 
4.7  ;  and  on  account  of  its  high  specific  gravity  it  received  the 
name  of  barytes,  or  heavy  spar. 

It  is  formed  as  a  heavy  white  precipitate  when  sulphuric  acid, 
or  a  soluble  sulphate,  is  added  to  a  solution  of  a  barium  salt.  It  is 
insoluble  in  water  and  only  very  slightly  soluble  in  dilute  acids. 


Barium  Sulphide  589 

It  is  soluble  in  hot  concentrated  sulphuric  acid,  especially  when 
freshly  precipitated  ;  and  the  solution  deposits,  on  cooling,  an  acid 
sulphate,  BaSO4,H2SO4.  On  exposure  to  moisture  the  solution 
deposits  crystals  of  BaSO4,H2SO4,2H2O.  Both  of  these  com- 
pounds, in  contact  with  water,  yield  insoluble  normal  barium 
sulphate  and  sulphuric  acid. 

Precipitated  barium  sulphate  is  largely  used  as  a  pigment, 
known  as  permanent  white. 

Barium  Nitrate,  Ba(NO3)2,  is  obtained  by  dissolving  the  native 
carbonate,  or  the  sulphide,  in  dilute  nitric  acid.  It  is  also  formed 
by  double  decomposition,  when  hot  saturated  solutions  of  sodium 
nitrate  and  barium  chloride  are  mixed.  The  salt  crystallises  in 
large  colourless  octahedrons.  100  parts  of  water  at  the  ordinary 
temperature  dissolve  9  parts,  and  at  100°,  32.2  parts  of  barium 
nitrate.  When  strongly  heated  it  is  converted  into  barium  oxide, 
with  the  evolution  of  nitrogen  peroxide,  oxygen,  and  nitrogen. 

Barium  nitrate  is  used  in  pyrotechny,  in  the  preparation  of 
mixtures  for  green  fire. 

Barium  Sulphide,  BaS,  is  obtained  by  methods  analogous  to 
those  for  preparing  calcium  sulphide  (page  583),  which  it  closely 
resembles  in  its  properties. 


CHAPTER  VII 
ELEMENTS  OF  GROUP  II.  (FAMILY  B.) 

Zinc,  Zn 65.4 

Cadmium,  Cd 112.4 

Mercury,  Hg         ......     200 

THE  three  elements  composing  this  family  do  not  exhibit  such 
a  close  resemblance  to  each  other  as  exists  between  barium, 
strontium,  and  calcium  ;  for  although  zinc  and  cadmium  are  very 
closely  related,  mercury  in  many  respects  differs  widely  from  these, 
and  from  all  the  other  elements  in  the  same  group. 

Cadmium  and  zinc  are  almost  invariably  found  associated 
together  in  nature,  they  are  both  fairly  permanent  in  the  air, 
and  both  readily  take  fire  and  burn  when  strongly  heated, 
forming  the  oxides.  Both  are  acted  upon  by  dilute  hydrochloric 
and  sulphuric  acids,  with  evolution  of  hydrogen,  and  most  of  their 
salts  are  isomorphous. 

Mercury  is  peculiar  in  being  liquid  at  ordinary  temperatures. 
Zinc  and  cadmium  melt  at  430°  and  320°  respectively,  while 
mercury  melts  at  —38.8°.  It  is  quite  unacted  upon  by  oxygen  at 
ordinary  temperatures,  and  combines  with  extreme  slowness  when 
heated.  Its  oxide,  also,  is  readily  decomposed  by  heat  into  its 
elements. 

Dilute  hydrochloric  and  sulphuric  acids  are  entirely  without 
action  upon  it,  and  it  forms  no  hydroxide. 

Mercury  also  differs  from  zinc  and  cadmium  in  forming  two 
elementary  ions,  giving  rise  to  mercurous  and  mercuric  salts. 
Both  zinc  and  cadmium  have  only  one  ion  and  form  only  one 
series  of  salts. 

The  hydroxide  of  zinc,  Zn(HO)2,  differs  from  the  corresponding 
cadmium  compound,  in  being  soluble  in  alkaline  hydroxides. 

These  three  elements  resemble  each  other,  and  differ  from 
those  of  family  A  of  this  group,  in  that  they  can  be  volatilised; 
mercury  at  a  temperature  about  357°,  cadmium  and  zinc  at 
temperatures  approaching  1000°. 

These  three  elements  are  also  alike,  in  that  their  vapours  con- 
sist of  mono-atomic  molecules. 

59° 


Zinc  591 

ZINC. 

Symbol,  Zn.     Atomic  weight =65.37. 

Oeeurrenee. — Zinc  is  stated  to  have  been  found  in  Australia  in 
the  uncombined  condition  ;  with  this  exception,  it  is  always  met 
with  in  combination,  chiefly  as  carbonate  in  calamine  or  zinc-spar, 
ZnCO3,  and  as  sulphide  in  zinc-blende,  or  black-jack,  ZnS.  Other 
ores  are  red  zinc  ore,  ZnO  ;  &s\&franklinite,  (ZnFe)O,Fe2O3. 

Gahnite,  or  zinc-spinnelle,  has  the  composition  ZnO,Al2O3. 

Modes  Of  Formation. — The  ores  chiefly  employed  for  the  pre- 
paration of  zinc  are  the  carbonate  and  sulphide,  although  in  New 
Jersey  the  red  oxide  and  franklinite  are  used.  The  process  con- 
sists of  two  operations,  namely,  first,  the  conversion  of  the  ore  into 
oxide  of  zinc,  by  calcination ;  and,  second,  the  reduction  of  the  oxide 
by  means  of  coal  at  a  high  temperature.  The  calcination  of  the 
natural  carbonate  is  readily  accomplished,  this  compound  merely 
giving  up  its  carbon  dioxide  at  a  high  temperature — 

ZnCO3  =  ZnO  +  CO2. 

In  the  case  of  zinc-blende,  the  operation  consists  in  the  oxida- 
tion of  both  the  sulphur  and  the  zinc  by  atmospheric  oxygen,  thus — 

ZnS  +  3O  =  ZnO  +  SO2. 

Considerable  care  has  to  be  exercised  in  order  to  prevent  the 
formation  of  zinc  sulphate,  which,  in  the  subsequent  operation, 
would  be  reconverted  into  sulphide,  and  so  lost.  The  finely 
crushed  calcined  ore  is  mixed  with  coke  or  coal  and  heated  to 
bright  redness  in  earthenware  retorts,  when  the  oxide  is  reduced, 
with  the  formation  of  carbon  monoxide,  and  the  metal  distils  and 
is  collected  in  iron  receivers.  Zinc  ores  frequently  contain  small 
quantities  of  cadmium,  and  as  this  metal  is  more  readily  volatilised 
than  zinc,  it  passes  over  in  the  first  portions  of  the  distilled 
product. 

The  two  processes  now  almost  exclusively  in  use  for  the  reduc- 
tion of  zinc,  known  as  the  Silesian  and  the  Belgian  process,* 
differ  only  in  metallurgical  details,  &c. 

*  The  old  method,  known  as  the  English  process,  or  distillation  per 
descensum,  is  entirely  obsolete.  For  details  of  this  and  all  other  metallurgical 
processes,  the  student  is  referred  to  treatises  on  metallurgy,  such  as  Percy. 


592  Inorganic  Chemistry 

Commercial  zinc  usually  contains  carbon,  iron,  and  lead,  and 
occasionally  arsenic  and  cadmium.  It  may  be  obtained  in  a  higher 
degree  of  purity  by  careful  distillation,  but  pure  zinc  is  best  ob- 
tained by  first  preparing  the  pure  carbonate  by  precipitation,  and 
then  calcining  and  finally  reducing  with  charcoal  obtained  from 
sugar. 

Properties. — Zinc  is  a  bluish-white,  highly  crystalline,  and 
brittle  metal.  At  a  temperature  of  300°  it  can  be  readily  powdered 
in  a  mortar,  while  between  100°  and  150°  it  admits  of  being  drawn 
into  wire  or  rolled  into  thin  sheet.  The  presence  of  a  small 
quantity  of  lead  greatly  enhances  this  property,  but  is  detrimental 
when  the  zinc  is  required  for  making  brass.  Zinc  which  has  been 
either  rolled  or  drawn  no  longer  becomes  brittle  when  cold,  but 
retains  its  malleability. 

Zinc  melts  at  a  temperature  about  430,°  and  when  heated  in  air 
much  beyond  this  point  the  metal  takes  fire  and  burns  with  a  bluish- 
white  flame,  the  brilliancy  of  which  becomes  dazzling  if  a  stream  of 
oxygen  be  projected  upon  the  burning  mass.  The  product  of  its 
combustion  is  zinc  oxide,  ZnO,  which  forms  a  soft,  white,  flocculent 
substance  resembling  wool,  and  formerly  known  as  philosophers 
wool. 

The  boiling-point  of  zinc  is  about  930°. 

Zinc  is  permanent  in  dry  air  at  ordinary  temperatures,  but  when 
exposed  to  moist  air  it  tarnishes  superficially  ;  it  is  also  unattacked 
by  water  at  the  boiling  temperature.  It  is  soluble  in  a  hot  solution 
of  sodium  or  potassium  hydroxide,  with  evolution  of  hydrogen 

(P-  175)- 

Pure  zinc  is  scarcely  acted  upon  by  pure  sulphuric  or  hydrochloric 
acid,  either  dilute  or  strong.  The  presence  of  small  quantities  of 
impurities,  however,  determines  the  solution  of  the  metal  with  the 
rapid  evolution  of  hydrogen,  hence  ordinary  commercial  zinc  is 
readily  attacked  by  these  acids,  and  also  decomposes  water  at  the 
boiling-point,  with  the  evolution  of  hydrogen.* 

*  The  difference  between  the  behaviour  of  acids  towards  pure  and  com- 
mercial zinc  was  formerly  explained  on  the  ground  that  the  impurities  present 
formed  with  the  zinc  a  voltaic  couple,  whereby  local  electric  currents  were  set 
up,  while  in  the  case  of  pure  zinc  no  such  action  took  place.  The  recent 
observations  of  Pullinger  (Chem.  Soc.,  57)  and  Weeren  (Berichte,  24)  show  that 
this  is  not  a  complete  explanation.  Weeren  concludes  that  the  insolubility  of 
pure  zinc  in  dilute  acids  is  due  to  the  formation  of  a  film  of  condensed  hydrogen 
upon  the  surface  of  the  metal,  which  stops  all  further  action.  The  addition  of 
oxidising  agents,  such  as  hydrogen  peroxide,  or  dilute  sulphuric  acid  which  has 


Zinc  Oxide  593 

Zinc  is  extensively  used  in  the  process  of  galvanising  iron,  which 
consists  in  coating  iron  with  a  film  of  zinc,  not  by  electrical  deposi- 
tion, as  would  be  implied  by  the  name,  but  by  dipping  the  iron 
into  a  bath  of  molten  zinc.  The  layer  of  zinc  preserves  the  iron 
from  rusting.  Galvanised  iron  is  better  able  to  withstand  the 
action  of  air  and  moisture  than  tinned  iron,  hence  it  is  extensively 
used  for  wire  netting,  corrugated  roofing,  water  tanks,  and  other 
purposes  where  the  metal  is  exposed  to  the  oxidising  influence  of 
air  and  water. 

Alloys  of  Zinc. — Zinc  forms  a  number  of  useful  alloys,  the  most 
important  of  which  are  the  various  forms  of  brass  (see  Copper). 
With  certain  metals,  such  as  tin,  copper,  and  antimony,  zinc  will 
mix  in  all  proportions  ;  while  with  others,  such  as  lead  and  bismuth, 
it  is  only  possible  to  obtain  solid  alloys  of  definite  composition. 
When,  therefore,  lead  and  zinc  are  melted  together,  although  in 
the  molten  condition  the  mixture  is  homogeneous,  on  cooling  the 
metals  separate  into  two  layers,  the  lighter  zinc  rising  to  the  surface. 
The  separation  of  the  metals,  however,  is  not  perfect,  for  the  zinc 
will  have  dissolved  a  certain  quantity  of  the  lead  (1.2  per  cent.), 
and  the  lower  layer  of  lead  is  found  to  have  dissolved  a  small 
proportion  of  zinc  (1.6  per  cent),  just  as  water  and  ether,  when 
shaken  together,  separate  into  two  layers,  the  uppermost  being  an 
ethereal  solution  of  water,  and  the  lower  an  aqueous  solution  of 
ether. 

This  property  is  made  use  of  in  the  extraction  of  silver  from  lead 
(see  p.  561). 

The  so-called  German  silver,  or  nickel  silver^  is  a  nearly  white 
alloy  of  copper,  nickel,  and  zinc. 

Bronze  coinage  consists  of  95  parts  of  copper,  4  of  tin,  and  i  of 
zinc,  the  small  proportion  of  zinc  giving  to  the  alloy  an  increased 
hardness  and  durability. 

Zinc  Oxide,  ZnO,  the  only  oxide  of  zinc,  occurs  native  as  red 
zinc  ore,  the  colour  being  due  to  the  presence  of  manganese.  It  is 

been  electrolysed,  and  therefore  contains  presulphuric  acid,  tends  to  destroy 
this  film  by  oxidising  the  hydrogen,  and  therefore  promotes  the  solution  of  the 
zinc.  He  also  finds,  that  by  mechanically  removing  this  layer  Of  hydrogen, 
either  by  constantly  brushing  the  metallic  surface  or  placing  the  materials 
under  reduced  pressure,  the  solution  of  the  zinc  by  the  acid  is  promoted.  It  is 
also  found  that  the  character  of  the  surface  of  the  metal,  whether  smooth  or 
rough,  affects  the  result ;  zinc  that  is  unacted  upon  when  its  surface  is  perfectly 
smooth  is  more  readily  attacked  by  the  dilute  acid  when  its  surface  is  rough. 

2  P 


594  Inorganic  Chemistry 

formed  as  a  soft  white  substance  when  zinc  is  burnt  in  the  air.  It 
is  manufactured  under  the  name  of  zinc  white  by  the  combustion 
of  zinc,  the  fumes  being  led  into  condensing-chambers,  where  the 
oxide  collects. 

Zinc  oxide  is  a  pure  white  substance,  which  when  heated  becomes 
yellow,  but  again  becomes  white  on  cooling.  When  strongly  heated 
in  oxygen,  it  may  be  obtained  in  the  form  of  hexagonal  crystals  ; 
such  crystals  are  occasionally  found  in  the  cooler  parts  of  zinc 
furnaces.  The  oxide  does  not  fuse  in  the  oxyhydrogen  flame,  but, 
like  lime,  under  these  circumstances  it  becomes  intensely  incan- 
descent ;  for  some  time  after  being  so  heated  it  appears  phos- 
phorescent in  the  dark.  It  is  insoluble  in  water,  and  does  not 
combine  directly  with  water  to  form  the  hydroxide.  It  dissolves 
in  acids,  giving  rise  to  the  different  zinc  salts.  Zinc  oxide  is  largely 
used  in  the  place  of"  white  lead"  as  a  pigment  ;  although  it  does 
not  equal  white  lead  in  covering  power,  or  body,  it  possesses  the 
advantage  of  not  being  blackened  by  exposure  to  atmospheric 
sulphuretted  hydrogen. 

Zinc  Hydroxide,  Zn(HO)2,  is  formed  as  a  white  flocculent  pre- 
cipitate, when  either  sodium  or  potassium  hydroxide,  or  a  solution 
of  ammonia,  is  added  to  a  solution  of  zinc  sulphate.  The  compound 
is  soluble  in  an  excess  of  either  alkali,  and  is  deposited  from  a 
strong  solution  in  regular  octahedra  of  the  hydrated  hydroxide, 
Zn(HO)2,H2O.  Both  of  these  compounds  on  heating  readily  lose 
water,  and  are  converted  into  the  oxide. 

Zinc  Chloride,  ZnCl2,  is  formed  by  the  direct  combination  of  zinc 
with  chlorine,  or  by  the  action  of  hydrochloric  acid  upon  the  metal. 
It  is  also  obtained  in  the  anhydrous  state  by  distilling  a  mixture  of 
mercuric  chloride  and  zinc,  or  a  mixture  of  anhydrous  zinc  sulphate 
and  calcium  chloride. 

It  is  usually  prepared  on  a  large  scale  by  dissolving  zinc  in 
hydrochloric  acid,  and  after  precipitating  any  manganese  and  iron, 
the  liquid  is  boiled  down  in  enamelled  iron  vessels,  until  on  cooling 
it  solidifies  ;  it  is  usually  cast  into  sticks. 

Zinc  chloride  is  a  soft,  white,  easily  fusible  solid,  which  volatilises 
and  distils  without  decomposition.  It  is  extremely  deliquescent, 
and  readily  soluble  in  water  and  in  alcohol,  its  solution  being 
powerfully  caustic.  From  a  strong  aqueous  solution  deliquescent 
crystals  are  deposited,  having  the  composition  ZnCl2,H2O. 

When  the  aqueous  solution  is  evaporated,  partial  decomposition 
takes  place,  hydrochloric  acid  being  evolved  and  basic  compounds 


Zinc  Sulphide  595 

being  precipitated,  consisting  of  combinations  of  the  chloride  and 
oxide.  Hence,  during  the  concentration  of  the  liquid  in  the  pre- 
paration of  zinc  chloride,  hydrochloric  acid  is  added  to  redissolve 
this  compound. 

A  paste  made  by  moistening  zinc  oxide  with  zinc  chloride  rapidly 
sets  to  a  hard  mass  ;  this  mixture,  under  the  name  of  oxychloride 
of  zinc,  is  employed  in  dentistry  as  a  filling  or  stopping  for  teeth. 

Zinc  chloride  unites  with  alkaline  chlorides,  forming  a  series  of 
crystalline  double  salts  having  the  general  formula  ZnCl2,2RCl. 

Zine  Sulphate,  ZnSO4,  is  formed  when  zinc  is  dissolved  in 
sulphuric  acid.  It  is  obtained  on  a  large  scale  by  roasting  the 
natural  sulphide,  whereby  it  is  partially  converted  into  the  sulphate, 
which  is  then  extracted  with  water. 

The  salt  crystallises  from  its  aqueous  solution  at  ordinary  tem- 
peratures in  colourless  rhombic  prisms,  ZnSO4,7H2O,  isomorphous 
with  MgSO4,7H2O.  It  is  extremely  soluble  in  water  :  100  parts  of 
water  at  the  ordinary  temperature  dissolve  160  parts,  and  at  100°, 
653.6  parts  of  the  crystalline  salts.  When  exposed  to  the  air,  the 
crystals  slowly  effloresce,  and  if  placed  in  vacuo  over  sulphuric 
acid,  or  if  heated  to  100°,  they  lose  six  molecules  of  water,  leaving 
the  monohydrated  salt  ZnSO4,H2O.  At  a  temperature  about  300° 
this  is  converted  into  the  anhydrous  compound,  and  at  a  white 
heat  it  gives  off  sulphur  dioxide  and  oxygen,  leaving  the  oxide. 

The  hydrated  salt,  ZnSO4,6H2O,  is  obtained  in  the  form  of  mono- 
symmetric  crystals,  when  the  salt  is  deposited  at  temperatures 
above  40°.  This  compound  is  isomorphous  with  MgSO4,6H2O. 

Zinc  sulphate  combines  with  alkaline  sulphates,  forming  a  series 
of  double  salts,  having  the  general  formula  ZnSO4,R2SO4,6H2O, 
which  are  also  isomorphous  with  the  corresponding  magnesium 
compounds  (page  576). 

Zinc  sulphate,  in  common  with  all  the  soluble  salts  of  zinc,  has 
an  astringent  taste,  and  is  poisonous. 

Zine  Sulphide,  ZnS. — The  natural  compound,  zinc-blende,  is 
usually  dark-brown  or  black,  and  exhibits  crystalline  forms  belong- 
ing to  the  regular  system.  The  mineral  ivurtzite  is  a  less  common 
variety  of  zinc  sulphide,  crystallising  in  hexagonal  prisms.  Zinc 
sulphide  is  obtained  as  a  white  amorphous  precipitate  when  an 
alkaline  sulphide  is  added  to  a  solution  of  a  zinc  salt,  or  when 
sulphuretted  hydrogen  is  passed  through  an  alkaline  solution  of  a 
zinc  salt. 

Precipitated  zinc  sulphide  is  insoluble  in  acetic  acid,  but  readily 


596  Inorganic  Chemistry 

dissolves  in  dilute  mineral  acids,  with  evolution  of  sulphuretted 
hydrogen  ;  hence  the  compound  is  not  formed  when  sulphuretted 
hydrogen  is  passed  through  a  solution  of  a  zinc  salt  containing  a 
free  mineral  acid. 

Zinc  Carbonate,  ZnCO3,  is  obtained  as  a  white  powder  when 
hydrogen  sodium  carbonate  is  added  to  a  solution  of  zinc  sulphate. 

If  normal  sodium  carbonate  be  employed,  the  precipitated  zinc 
compound  consists  of  a  basic  carbonate,  whose  composition  varies 
with  the  conditions  of  temperature  and  concentration  of  the  liquids. 

A  basic  carbonate, having  the  composition  ZnCO3,2Zn(H O)2, H2O, 
is  employed  as  a  pharmaceutical  preparation  under  the  name  zinci 
carbonas. 

CADMIUM. 

Symbol,  Cd.     Atomic  weight  =  112. 4. 

Occurrence. — Cadmium  is  never  found  in  the  uncombined  state. 
The  only  natural  compound  of  which  cadmium  is  the  chief  con- 
stituent is  the  extremely  rare  mineral  greenockite,  which  is  the 
sulphide,  CdS.  Cadmium  occurs  in  small  quantities  in  many  zinc 
ores,  such  as  the  sulphide  and  carbonate  ;  and  in  the  process  of 
extracting  zinc  from  these  ores,  the  cadmium  is  obtained  in  the 
first  portions  of  the  product  of  the  distillation,  partly  as  metal 
and  partly  as  oxide. 

Mode  Of  Formation.— The  crude  product  of  distillation  is  dis- 
solved in  dilute  sulphuric  or  hydrochloric  acid,  and  the  cadmium 
precipitated  as  sulphide  by  means  of  sulphuretted  hydrogen.  The 
cadmium  sulphide  is  then  dissolved  in  strong  hydrochloric  acid, 
and  precipitated  as  carbonate  by  means  of  ammonium  carbonate. 
The  washed  and  dried  carbonate  is  first  converted  into  oxide  by 
calcination,  and  finally  mixed  with  charcoal  and  distilled. 

Properties. — Cadmium  is  a  bluish- white  metal  resembling  zinc 
in  appearance,  but  much  more  malleable  and  ductile.  It  tarnishes 
superficially  on  exposure  to  the  air,  and,  when  strongly  heated, 
burns  with  the  formation  of  a  brown  smoke  of  cadmium  oxide, 
CdO.  The  metal  is  attacked  by  dilute  hydrochloric  and  sulphuric 
acids,  with  the  evolution  of  hydrogen.  It  readily  dissolves  in  nitric 
acid,  yielding  the  nitrate,  with  the  formation  of  oxides  of  nitrogen. 
Cadmium  is  less  electro-positive  than  zinc,  and  is  precipitated  in 
the  metallic  condition  from  its  solutions  by  that  metal. 

Cadmium  melts  at  320°,  and  boils  about  745°.     When  volatilised 


Mercury  597 

in  an  atmosphere  of  hydrogen,  it  forms  crystals  belonging  to  the 
regular  system. 

Cadmium  Oxide,  CdO,  is  formed  as  a  brown  fume  or  smoke 
when  cadmium  burns  in  the  air.  It  may  be  obtained  by  heating 
the  carbonate  or  nitrate.  That  obtained  by  the  ignition  of  the 
latter  salt  is  in  the  form  of  minute  crystals,  having  a  bluish-black 
appearance.  Cadmium  oxide  is  insoluble  in  water,  but  dissolves 
in  acids  yielding  cadmium  salts.  It  is  infusible  in  the  oxyhydrogen 
flame,  but  is  readily  reduced  when  heated  on  charcoal  before  the 
blowpipe  ;  and  the  reduced  metal,  as  it  volatilises  and  burns,  forms 
a  characteristic  brown  incrustation  of  oxide  upon  the  charcoal. 

Cadmium  Chloride,  CdCl2,  is  obtained  by  the  action  of  hydro- 
chloric acid  upon  the  metal  or  the  oxide.  The  salt  is  deposited 
from  the  solution  in  white  silky  crystals,  having  the  composition 
CdCl2,2H2O.  On  exposure  to  the  air  the  crystals  effloresce,  and 
when  heated  become  anhydrous. 

Cadmium  Sulphide,  CdS,  is  obtained  as  a  bright  yellow  preci- 
pitate when  sulphuretted  hydrogen  is  passed  through  a  solution  of 
a  cadmium  salt.  The  precipitate  is  soluble  in  concentrated  hydro- 
chloric and  nitric  acids,  and  in  warm  dilute  sulphuric  acid.  Cad- 
mium sulphide  is  insoluble  in  ammonium  sulphide  ;  thjs  property 
readily  distinguishes  it  from  arsenious  sulphide,  which  in  colour 
it  closely  resembles. 

Cadmium  sulphide  is  used  as  a  pigment,  both  in  oil  and  water- 
colours. 


MERCURY. 

Symbol,  Hg.     Atomic  weight  =  200. 

Occurrence. — In  the  uncombined  state  mercury  is  met  with  in 
small  globules,  disseminated  through  its  ores,  especially  the  sul- 
phide. It  is  also  occasionally  found  as  an  amalgam  with  silver 
and  gold.  The  principal  ore  is  cinnabar,  HgS,  and  the  chief 
mines  of  this  ore  are  those  of  Almaden  (Spain),  Idria  (Carniola), 
California,  and  the  Bavarian  Palatinate. 

Modes  Of  Formation. — Mercury  may  be  obtained  from  the 
natural  sulphides  by  either  roasting  the  ore,  whereby  the  sulphur 
is  oxidised  to  sulphur  dioxide  and  the  metal  liberated,  or  by  dis- 
tillation in  closed  retorts  with  lime,  when  calcium  sulphide  and 
sulphate  are  formed,  and  the  mercury  set  free.  The  first  method 
is  almost  exclusively  employed. 


598 


Tnorganic  Chemistry 


At  Idria  the  crude  ore,  consisting  of  cinnabar  mixed  with  shale 
and  earthy  matters,  is  roasted  in  a  furnace,  upon  perforated  arches, 
/z,  n' ;  p,  p'y  Fig.  143.  The  action  of  the  fire  and  heated  air  is  to 
oxidise  the  sulphur  and  volatilise  the  mercury,  and  the  gases  and 
vapours  together  pass  through  a  series  of  flues  or  chambers,  C,  C, 
where  the  mercury  condenses. 

By  the  use  of  a  reverberatory  furnace  (the  Alberti  furnace),  the 
process  can  be  made  continuous  The  ore  is  fed  into  the  furnace 


FIG.  143. 


through  a  hopper,  and  the  calcined  residue  is  raked  out  through 
an  opening  at  the  opposite  end  of  the  hearth.  The  gases  are 
passed  first  through  iron  pipes  kept  cool  by  water,  and  then 
through  a  series  of  chambers  where  the  remaining  metal  is 
condensed. 

The  method  adopted  at  Almaden  is  essentially  the  same  as  the 

Idrian   process,   except    that    the 
condensation    takes    place    in    a 
series    of    pear-shaped    earthen- 
FIG.  144.  ware  vessels,  called  aludels,  which 

are  connected  together  as  shown 

in  Fig.  144.  Usually  six  rows  of  forty-seven  such  aludels  are 
connected  with  six  openings  in  a  chamber  immediately  above  the 
furnace. 

The  impure  mercury  is  freed  from  mechanically  mixed  impurities 
by  straining  or  filtering  through  chamois  leather,  but  from  metals 
in  solution,  such  as  zinc,  tin,  lead,  and  others,  it  is  purified  by 
distillation.  For  laboratory  purposes,  pure  mercury  is  best  ob- 
tained by  distillation  in  vacuo,  by  means  of  the  apparatus  shown  in 


Mercury 


599 


Fig.  145  (Clarke).  In  this  arrangement  the  mercury  is  distilled 
in  a  Sprengel  vacuum.  The  mercury  (previously  cleaned  by  being 
thoroughly  agitated  with  mercuric  nitrate)  is  placed  in  the  reser- 
voir R,  which  is  then  placed  upon  the  upper  shelf  S,  and  by  means 
of  the  clamp,  mercury  is  allowed  to  pass  into  the  long  wide  tube  T, 
and  up  into  the  bulb.  The  air  in  the  tube  and  bulb  escapes  down 
the  narrow  inner  tube,  which  reaches  nearly  to  the  top  of  the  bulb> 


FIG.  145. 


as  seen  in  the  enlarged  detail,  t.  The  mercury  is  allowed  to  rise 
in  the  bulb  and  fall  down  the  long  inner  tube,  after  the  manner  of 
the  Sprengel  pump.  The  reservoir  is  then  placed  upon  the  lower 
adjustable  stand,  and  its  height  so  arranged  that  the  mercury  in 
the  bulb  falls  to  the  position  shown  in  the  figure.  This  space  is  a 
Torricellian  vacuum.  The  mercury  is  then  heated  by  a  ring- 
burner  B,  and  the  whole  is  protected  from  draught  by  the  hood  h. 


600  Inorganic  Chemistry 

As  the  mercury  distils,  it  passes  down  the  inner  tube,  and  by  its 
fall  continues  to  preserve  the  Sprengel  vacuum  within  the  bulb. 

Properties. — At  ordinary  temperatures  mercury  is  a  bright,  silver- 
white  liquid  metal  (hence  its  old  name  quicksilver,  i.e.  live  silver). 
When  cooled  to  —  38.8°  it  solidifies  to  a  highly  crystalline  solid, 
which  is  ductile  and  malleable,  and  softer  than  lead.  When  the 
liquid  is  cooled  it  contracts  uniformly  until  the  solidifying  point 
is  reached,  when  considerable  contraction  takes  place.  Solid 
mercury,  therefore,  is  denser  than  the  liquid  metal,  and  sinks  in 
it.  The  specific  gravity  of  liquid  mercury  at  o°  is  13.596,  while 
that  of  the  solid  at  its  melting-point  is  14.193.  Mercury  in 
extremely  thin  films  appears  a  violet  colour  by  transmitted  light. 

Under  a  pressure  of  760  mm.  mercury  boils  at  357«25°>  givmg 
a  colourless  vapour.  The  density  of  mercury  vapour  referred  to 
hydrogen  is  100.15  J  hence  this  element,  like  its  associates  in  the 
family  to  which  it  belongs,  consists  of  mono-atomic  molecules  when 
in  a  state  of  vapour.  Mercury  gives  off  vapour  even  at  ordinary 
temperatures,  and  a  gold  leaf  suspended  over  mercury  in  a  stop- 
pered bottle  gradually  becomes  white  upon  the  surface,  owing  to 
its  amalgamation  with  the  mercurial  vapour. 

The  vapour  of  mercury  is  poisonous,  giving  rise  to  salivation. 

Mercury  does  not  tarnish  on  exposure  to  the  air,  and  is  unacted 
upon  by  a  large  number  of  gases  ;  hence  this  liquid  is  invaluable 
to  the  chemist,  affording  a  means  of  collecting  and  measuring 
gases  which  are  soluble  in  water. 

When  submitted  to  prolonged  heating  in  the  air  it  is  slowly 
converted  into  the  red  oxide,  which  at  a  higher  temperature  is 
again  decomposed  into  its  elements. 

Mercury  is  obtained  in  the  form  of  a  dull-grey  powder  when  it 
is  shaken  up  with  oil  or  triturated  with  sugar,  chalk,  or  lard.  This 
operation  is  known  as  deadening,  and  is  made  use  of  in  the  pre- 
paration of  mercurial  ointment.  The  grey  powder  consists  simply 
of  very  finely  divided  mercury  in  the  form  of  minute  globules. 

Mercury  is  not  attacked  by  hydrochloric  acid.  Strong  sulphuric 
acid  is  without  action  upon  it  in  the  cold,  but  when  heated  the 
metal  dissolves,  with  evolution  of  sulphur  dioxide.  Strong  nitric 
acid  rapidly  attacks  it,  with  formation  of  mercuric  nitrate  and 
oxides  of  nitrogen.  Cold  dilute  nitric  acid  slowly  dissolves  it, 
forming  mercurous  nitrate. 

Alloys  Of  Mercury. — When  mercury  is  one  of  the  constituents 
of  an  alloy  the  mixture  is  called  an  amalgam.  Most  metals  will 


Salts  of  Mercury  60 1 

form  an  amalgam  with  mercury.  In  some  cases,  as  with  the 
alkali  metals,  the  union  is  attended  with  great  rise  of  temperature. 
In  other  cases,  as  with  tin,  an  absorption  of  heat  takes  place. 

Sodium  and  potassium  amalgams  are  obtained  by  dissolving 
various  amounts  of  the  metals  in  mercury.  In  contact  with  water 
they  are  decomposed,  hydrogen  being  evolved  and  the  alkaline 
hydroxide  formed.  On  this  account  sodium  amalgam  is  frequently 
used  in  the  laboratory  as  a  reducing  agent.  When  heated  to 
440°  these  amalgams  leave  behind  crystalline  compounds,  K2Hg 
and  Na3Hg,  which  spontaneously  inflame  in  contact  with  the  air. 

Zinc  amalgams  are  only  very  slowly  acted  upon  by  dilute  sul- 
phuric acid  ;  therefore,  by  the  superficial  amalgamation  of  the 
zinc  plates  used  for  galvanic  batteries,  the  same  result  is  ob- 
tained as  though  the  zinc  were  perfectly  pure  (see  page  592),  and 
no  solution  of  zinc  takes  place  until  the  electric  circuit  is  closed. 

Tin  amalgams  are  employed  for  the  construction  of  ordinary 
mirrors. 

Amalgams  of  gold,  and  also  copper  and  zinc,  are  used  in 
dentistry  as  a  filling  or  stopping  for  teeth. 

Oxides  of  Mercury. — Two  oxides  are  known,  namely,  mercu- 
rous  oxide,  Hg2O,  and  mercuric  oxide,  HgO. 

Mereurous  Oxide,  Hg2O,  is  obtained  as  an  unstable  dark-brown 
or  black  powder  when  sodium  hydroxide  is  added  to  mercurous 
chloride.  When  exposed  to  the  light,  or  when  gently  heated,  it  is 
converted  into  mercuric  oxide  and  mercury. 

Mereurie  Oxide,  HgO,  is  produced  in  small  quantity  by  the  pro- 
longed heating  of  mercury  in  contact  with  air,  or  by  igniting  the 
nitrate.  It  is  prepared  on  a  large  scale  by  heating  an  intimate 
mixture  of  mercuric  nitrate  and  mercury.  Obtained  by  these 
methods,  it  is  a  brick- red  crystalline  powder  ;  but  when  sodium 
hydroxide  is  added  to  a  solution  of  a  mercuric  salt,  the  oxide  is  pre- 
cipitated as  an  orange-yellow  amorphous  powder.  When  heated, 
mercuric  oxide  first  darkens  in  colour,  and  gradually  becomes 
almost  black,  but  returns  to  its  original  bright  red  colour  on  cool- 
ing. At  a  red  heat  it  is  completely  decomposed  into  its  elements. 

Salts  Of  Mercury.— Two  series  of  salts,  corresponding  to  the 
two  oxides,  are  known — (a)  mercurous  salts,  in  which  two  atoms  of 
the  hydrogen  of  the  acids  are  replaced  by  the  divalent  radical 
or  double  atom  (Hg2) ;  and  (/5)  mercuric  salts,  in  which  the  same 
amount  of  hydrogen  is  replaced  by  the  single  divalent  atom  (Hg). 
All  the  mercury  salts  are  poisonous. 


602  Inorganic  Chemistry 


(a)  MERCUROUS  SALTS. 

Mereurous  Chloride,  Hg2Cl2  (calomel},  is  met  with  in  small 
quantities  as  the  mineral  horn  mercury.  It  may  be  obtained  by 
the  addition  of  sodium  chloride  or  hydrochloric  acid  to  a  solution 
of  mercurous  nitrate.  On  a  large  scale  it  is  usually  prepared  by 
heating  a  mixture  of  mercuric  chloride  and  mercury,  when  the 
mercurous  chloride  sublimes  as  a  white  or  translucent  fibrous 
cake. 

When  a  mixture  of  mercuric  sulphate,  common  salt,  and  mercury 
is  heated,  mercurous  chloride  is  also  obtained,  thus — 

HgSO4  +  2NaCl  +  Hg  =  Na2SO4+Hg2Cl2. 

Calomel  is  perfectly  tasteless,  and  is  insoluble  in  water.  When 
heated  it  vaporises  without  fusing.  The  density  of  the  vapour 
that  is  formed  by  heating  mercurous  chloride  is  117.87,  which  is 
half  that  demanded  by  the  formula  Hg2Cl2.  It  has  been  shown, 
however,  that  the  compound  dissociates  when  vaporised  into 
mercuric  chloride  and  mercury.*  Boiling  hydrochloric  acid  de- 
composes mercurous  chloride  into  mercury,  which  separates  out, 
and  mercuric  chloride,  which  dissolves. 

Mereurous  Nitrate,  Hg2(NO3)2,  is  deposite'd  in  the  form  of 
colourless  monosymmetric  crystals  containing  2H2O,  from  a  solu- 
tion of  mercury  in  cold  dilute  nitric  acid.  The  salt  is  soluble  in 
water  acidulated  with  nitric  acid,  but  an  excess  of  water  causes 
the  precipitation  of  a  basic  nitrate  having  the  composition — 

Hg2(N03)2,Hg20,H20  (or  2Hg2(NO3)(HO)), 

which,  on  boiling,  is  converted  into  mercuric  nitrate  and  mercury. 
If  either  this  or  the  normal  salt  be  boiled  in  the  presence  of  an 
excess  of  mercury,  a  basic  nitrate  of  the  composition — 

3Hg2(N03)2,2Hg20,2H20  (or  Hg2(NO3)2,4Hg2(NO3)(HO)), 

is  obtained. 

Mereurous  Sulphate,  Hg2SO4,  is  obtained  as  a  white  crystalline 
precipitate  when  dilute  sulphuric  acid  is  added  to  a  solution  of 
mercurous  nitrate.  It  is  very  slightly  soluble  in  water. 

*  Harris  and  Meyer,  Berichte,  June  1894. 


Mercuric  Iodide  603 


(j8)  MERCURIC  SALTS. 

Mercuric  Chloride,  HgC\2  (corrosive  sublimate),  is  formed  when 
chlorine  is  passed  over  heated  mercury.  It  is  prepared  on  a  large 
scale  by  heating  a  mixture  of  mercuric  sulphate  and  common  salt, 
a  small  quantity  of  manganese  dioxide  being  added  to  prevent, 
as  far  as  possible,  the  formation  of  mercurous  chloride.  The 
mercuric  chloride  sublimes  as  a  white  translucent  mass.  It  dis- 
solves in  water  to  the  extent  of  6.57  parts  in  100  parts  of  water  at 
10°,  and  54  parts  in  the  same  volume  of  water  at  100°,  forming  an 
acid  solution  from  which  the  salt  is  deposited  in  long  white  silky 
needles.  It  readily  melts,  and  volatilises  unchanged.  It  dissolves 
without  decomposition  in  nitric  acid  and  in  sulphuric  acid,  and 
volatilises  unchanged  from  its  solution  in  the  latter  acid  on 
boiling. 

Mercuric  chloride  is  a  violent  poison  :  the  best  antidote  is  albu- 
men, with  which  it  forms  an  insoluble  compound.  It  has  also 
strong  antiseptic  properties,  and  on  this  account  is  largely  used  by 
taxidermists. 

With  hydrochloric  acid,  mercuric  chloride  forms  two  crystalline 
double  chlorides,  HgCl2,HCl  and  2HgCl2,HCl;  and  with  the 
alkaline  chlorides  it  forms  a  number  of  similar  double  salts,  of 
which  the  ammonium  compound,  HgCl2,2NH4Cl,H2O,  was  known 
to  the  early  chemists  under  the  name  sal  alembroth. 

Mercuric  Iodide,  HgI2. — When  mercury  and  iodine  are  rubbed 
together  in  a  mortar,  and  moistened  with  a  small  quantity  of 
alcohol,  the  red  mercuric  iodide  is  formed.  It  is  also  obtained  by 
precipitation  from  a  solution  of  mercuric  chloride,  upon  the  addition 
of  potassium  iodide.  The  precipitate  first  appears  yellow,  but  in  a 
few  seconds  becomes  scarlet. 

Mercuric  iodide  is  insoluble  in  water,  but  readily  dissolves  in 
either  mercuric  chloride  or  potassium  iodide,  and  also  in  alcohol 
and  in  nitric  acid.  From  its  solutions  it  is  deposited  in  scarlet 
tetragonal  pyramids. 

Mercuric  iodide  is  dimorphous  ;  when  heated  to  about  150°  the 
scarlet  crystals  are  changed  into  bright  yellow  orthorhombic 
prisms.  At  ordinary  temperatures  this  yellow  form  is  unstable, 
and  on  being  lightly  touched  it  is  at  once  retransformed  into 
the  red  modification.  At  very  low  temperatures,  however,  the 
yellow  variety  is  the  more  stable  :  thus,  when  the  red  crystals 


604  Inorganic  Chemistry 

are  exposed  to  the  temperature  of  evaporating  liquid  oxygen,  they 
pass  into  the  yellow  variety. 

Mercuric  Nitrate,  Hg(NO3)2,  is  prepared  by  boiling  nitric  acid 
with  mercury,  until  sodium  chloride  produces  no  precipitate  with  a 
sample  of  the  liquid.  If  this  solution  be  evaporated  over  sulphuric 
acid,  deliquescent  crystals  are  obtained  of  2Hg(NO3)2,H2O,  while 
the  mother-liquor  has  the  composition  Hg(NO3)2,2H2O. 

Mercuric  nitrate  exhibits  a  great  tendency  to  form  basic  salts  : 
thus,  when  this  mother-liquor  is  boiled,  the  compound  Hg(NO3)2, 
HgO,2H2O  is  precipitated.  When  this  compound,  or  the  normal 
nitrate,  is  treated  with  an  excess  of  cold  water,  there  is  formed  the 
still  more  basic  salt  Hg(NO3)2,2HgO,H2O. 

Mercuric  Sulphide,  HgS  (#»»rf&zr).--When  mercury  and 
sulphur  are  triturated  together  in  a  mortar,  or  when  excess  of 
sulphuretted  hydrogen  is  passed  into  a  solution  of  a  mercuric  salt, 
mercuric  sulphide  is  obtained  as  a  black  amorphous  powder.  If 
this  be  sublimed,  it  is  obtained  as  a  red  crystalline  substance. 

Mercuric  sulphide  in  the  red  condition  is  also  obtained  by 
digesting  the  black  amorphous  product  for  some  hours  in  alkaline 
sulphides.  A  soluble  double  sulphide  is  first  formed,  which  when 
heated  is  decomposed,  with  the  deposition  of  red  mercuric  sulphide. 
This  compound  is  manufactured  on  a  large  scale  for  use  as  the 
pigment  vermilion. 

Mercuric  sulphide  is  insoluble  in  either  nitric,  hydrochloric,  or 
sulphuric  acid.  In  the  presence  of  an  alkali  it  is  soluble  in  sodium 
or  potassium  sulphide,  and  deposits  crystals  from  these  solutions 
having  the  composition  HgS,Na2S,8H2O,  and  HgS,K2S,5H2O 
respectively. 

Ammoniaeal  Mercury  Compounds.— These  may  be  regarded 
as  ammonium  salts,  in  which  two  atoms  of  hydrogen  in  ammonium 
(NH4)  have  been  replaced  by  either  (Hg2)  in  the  mercurous,  or  by 
(Hg)  in  the  mercuric  compounds  ;  the  two  atoms  so  replaced  being 
either  drawn  from  one  and  the  same  ammonium  group,  or  from  two. 


(a)  MERCUROUS  COMPOUNDS. 

Mercurous  Ammonium  Chloride,  (NH2Hg2)Cl,  is  the  black 
powder  produced  by  the  action  of  aqueous  ammonia  upon  calomel, 
thus— 

=  (NH2Hg2) 


Mercuric  Compounds  605 

Mercurous  Ammonium  Nitrate,  (NH2Hg2)NO3,  is  formed,  to- 
gether with  other  compounds,  when  aqueous  ammonia  is  added  to 
mercurous  nitrate. 

Mereurous  Diammonium  Chloride,  ^H3C1  (  H§"2  or  ^H^ 
HgaClo,  is  obtained  when  calomel  absorbs  dry  gaseous  ammonia. 
On  exposure  to  the  air  it  gives  up  its  ammonia,  and  is  reconverted 
into  mercurous  chloride. 


(j8)  MERCURIC  COMPOUNDS. 

Mercuric  Ammonium  Chloride,  (NH2Hg)Cl  (infusible  white 
precipitate],  is  formed  when  ammonia  is  added  to  a  solution  of 
mercuric  chloride  — 


Dimereurie  Ammonium  Chloride,  (NHg2)Cl,  is  obtained  by 
the  action  of  water  on  the  preceding  compound. 

Mercuric  Diammonium  Chloride,  ^^c!  }  Hg,or(N  H3)2HgCl2 

{fusible  white  precipitate},  is  obtained  by  adding  mercuric  chloride 
to  a  boiling  aqueous  solution  of  ammonium  chloride  and  ammonia, 
until  the  precipitate  which  first  forms  no  longer  dissolves.  On 
cooling,  the  solution  deposits  small  crystals  belonging  to  the 
regular  system. 

Oxy-dimereurie  Ammonium  Iodide,  (NH2Hg)I,HgO,  is  pro- 
duced by  the  action  of  aqueous  ammonia  upcn  mercuric  iodide, 
thus  — 

4NH3  +  2HgI2  +  H2O  =  (NH2Hg)I,HgO  +  3NH4I. 

It  is  readily  produced  as  a  brown  precipitate  by  adding  ammonia 
to  a  solution  of  mercuric  iodide  in  potassium  iodide  containing  an 
excess  of  potassium  hydroxide. 

The  alkaline  solution  of  potassium  mercuric  iodide  is  known  as 
NessleSs  solution,  and  constitutes  a  delicate  reagent  for  detecting 
the  presence  of  ammonia.  Minute  traces  of  free  ammonia  in  solu- 
tion produce  a  yellow  or  brown  coloration  with  this  test. 


CHAPTER   VIII 
THE  ELEMENTS  OF  GROUP  III 

Family  A.  Family  B. 


Scandium,  Sc      .  .  44.1 

Yttrium,  Y          .  .         89 

Lanthanum,  La .  .  139 

Ytterbizim,  Yb    .  .  172 


Boron,  B   .         .  .         n 

Aluminium,  Al  .  .         27.1 

Gallium,  Ga      .  .         70 

Indium,  In         .  .  115 

Thallium,  Tl     .  .  204 


WITH  the  exception  of  boron,  aluminium,  and  thallium,  the  mem- 
bers of  this  group  are  amongst  the  rarest  of  the  elements.*  Some 
of  these  occur  only  in  minute  traces  in  certain  ores  of  other  metals  : 
such  is  the  case  with  the  elements  gallium  and  indium,  which  are 
met  with  in  certain  specimens  of  zinc-blende,  the  ore  being  con- 
sidered rich  in  gallium  if  it  contains  as  much  as  0.002  per  cent,  of 
this  element.  Both  gallium  and  indium  were  discovered  by  means 
of  the  spectroscope  ;  the  latter  by  Reich  and  Richter  (1863),  and 
named  indium  on  account  of  two  characteristic  lines  in  the  indigo- 
blue  part  of  the  spectrum  ;  gallium  by  Lecocq  de  Boisbaudran 
(1875),  and  named  after  his  own  country.  The  spectrum  of  this 
metal  is  characterised  by  two  violet  lines.  One  of  the  most 
remarkable  properties  of  gallium  is  its  extremely  low  fusing-point, 
the  metal  melting  at  30.15°.  (For  a  comparison  of  the  properties 
of  gallium  with  Mendelejeff's  eka-aluminium^  see  page  124.) 

Others  of  these  elements  are  met  with  in  certain  rare  minerals, 
thus,  lanthanum  occurs  in  the  mineral  orthite  (from  Greenland)  ; 
and  both  yttrium  and  lanthanum  (associated  also  with  the  rare 
elements  cerium  and  erbium)  are  found  in  gadolinite  or  ytterbite 
(from  Ytterby). 

Boron  (the  typical  element  of  the  group)  is  the  only  non-metal  : 
all  the  others  exhibit  well-marked  metallic  properties.  They  all 
yield  sesquioxides  of  the  type  R2O3 ;  in  the  case  of  boron  this 
oxide,  B2O3,  is  acidic. 

*  For  detailed  descriptions  of  the  rare  elements,  the  student  is  referred  to 

larger  treatises,  or  to  chemical  dictionaries. 

606 


Boron  607 

Thallium  in  many  respects  is  peculiar.  It  forms  two  series  of 
compounds ;  in  one  class  it  functions  as  a  monovalent,  and  in  the 
other  as  a  trivalent  element.  In  some  of  its  properties  it  exhibits  a 
close  analogy  to  the  alkali  metals  ;  thus,  it  forms  a  soluble  strongly 
alkaline  hydroxide,  T1HO,  corresponding  to  KHO.  And  many  of 
its  salts,  such  as  the  sulphate,  T12SO4 ;  perchlorate,  T1C1O4,  and 
the  phosphates,  are  isomorphous  with  the  corresponding  potassium 
compounds. 

Thallium  also  shows  many  properties  in  common  with  lead, 
which  in  the  periodic  system  is  the  next  element  in  the  series 
(the  fourth  long  series).  Thus,  the  chloride,  like  lead  chloride, 
is  thrown  down  as  a  white  curdy  precipitate  on  the  addition  of 
hydrochloric  acid  to  a  soluble  salt  of  the  metal,  and  like  lead 
chloride,  thallous  chloride  is  soluble  in  hot  water.  Thallous 
iodide  also  closely  resembles  lead  iodide,  being  formed  as  a  yellow 
crystalline  precipitate  when  potassium  iodide  is  added  to  a  soluble 
thallous  salt. 

Metallic  thallium  also  bears  the  closest  resemblance  to  metallic 
lead. 

In  the  thallic  compounds  this  element  is  more  closely  related  to 
the  other  members  of  this  family :  thus,  thallic  oxide,  T12O3  ;  thallic 
chloride,  T1C13  ;  and  thallic  sulphide,  T12S3,  are  analogous  to  the 
corresponding  boron  compounds,  B2O3,  BC13,  B2S3. 


BORON. 

Symbol,  B.     Atomic  weight=n. 

Occurrence.—  The  element  boron  has  never  been  found  in  the 
free  state.  In  combination  it  occurs  principally  as  boric  acid  in 
volcanic  steam,  and  as  metallic  borates,  of  which  the  commonest 
are  tincal,  a  crude  sodium  borate,  or  borax,  Na2B4O7,10H2O  j 
boracite  and  colemanite,  or  borate  spar,  Ca2B6On  ;  and  boronatro- 
calcite,  or  ulexite,  Ca2B6O11,Na2B4O7,16H2O. 

Modes  of  Formation. — (i.)  Boron  may  be  prepared  by  heating 
boron  trioxide  with  either  sodium  or  potassium  in  a  covered 
crucible — 

2B2O3  +  6Na  =  3Na2O2  +  4B. 

The  fused  mass  is  boiled  with  dilute  hydrochloric  acid,  and  the 


608  Inorganic  Chemistry 

boron,  which  is  in  the  form  of  a  dark-brown  powder,  is  separated 
by  filtration. 

(2.)  The  element  may  also  be  obtained  by  heating  potassium 
borofluoride  with  potassium  — 

BF3,KF  +  3K  =  4KF  +  B. 

(3.)  Boron  is  also  formed  when  potassium  is  heated  in  the 
vapour  of  boron  trichloride  — 

BC13  +  3K  =  3KC1  +  B. 

Properties.  —  Boron,  as  obtained  by  these  methods,  is  a  dark 
greenish-brown  powder.  When  strongly  heated  in  air  it  burns, 
uniting  both  with  oxygen  and  nitrogen,  forming  a  mixture  of  boron 
trioxide,  B2O3,  and  boron  nitride,  BN.  It  is  unacted  upon  by  air 
at  ordinary  temperatures. 

Boron  has  no  action  upon  boiling  water,  but  cold  nitric  acid 
converts  it  into  boric  acid  — 

B  +  3HN03=H3BO3  +  3NO2. 
When  heated  with  sulphuric  acid  it  is  similarly  oxidised  — 


When  fused  with  alkaline  carbonates,  nitrates,  sulphates,  and 
hydroxides  it  forms  borates  of  the  alkali  metals,  thus  — 


2B  +  3Na3CO3  =  2Na3BO3 

2B  +  6KHO      =2K3BO3  +3H2. 

Boron  dissolves  in  molten  aluminium,  which  on  cooling  deposits 
crystals  of  a  compound  of  aluminium  and  boron.* 

Boron  Trioxide,  B2O3,  is  formed  when  boron  burns  in  the  air  or 
in  oxygen.  The  readiest  method  for  its  preparation  consists  in 
heating  boric  acid  to  redness,  when  it  fuses  and  gives  up  water  — 

2B(HO)3  =  3H2O  +  B2O3. 

Properties.  —  The  fused  mass  solidifies  to  a  transparent,  colour- 

*  This  compound  was  at  one  time  mistaken  for  an  allotropic  modification  of 
boron. 


Orthoboric  Acid  609 

less,  vitreous  solid,  which  gradually  absorbs  atmospheric  moisture 
and  becomes  opaque.  It  is  not  volatile  below  a  white  heat,  and 
on  this  account,  although  only  a  feeble  acid,  it  is  capable  at  high 
temperatures  of  displacing  strong  acids  which  are  volatile  from 
their  combinations  ;  thus,  when  boron  trioxide  is  fused  with  potas- 
sium sulphate,  potassium  borate  is  formed  and  sulphur  trioxide 
expelled — 

B2O3+3K2SO4  =  2B(KO)3  +  3SO3. 

Boron  trioxide  at  a  high  temperature  is  capable  of  dissolving 
many  metallic  oxides,  some  of  which  impart  to  the  fused  mass  a 
characteristic  colour. 

Boron  forms  three  oxyacids,  namely — 

Orthoboric  acid,  B(HO)3,  or  H3BO3. 

Metaboric  acid,  B2O2(HO)2,  or  H2B2O4,  or  B2O3,H2O. 

Pyroboric  acid,  B4O5(HO)2,  or  H2B4O7,  or  2B2O3,H2O. 

Orthoborie  Acid,  or  Boric  Acid,  B(HO)3,  occurs  naturally, 
both  in  the  waters  and  in  the  jets  of  steam  which  issue  from  the 
ground  in  many  volcanic  districts,  notably  in  Tuscany. 

The  actual  amount  of  boric  acid  in  these  natural  jets  of  steam 
or  soffioni  is  very  small  ;  but  as  the  steam  becomes  condensed  in 
the  pools  of  water  or  lagoons  which  often  surround  the  jets,  the 
amount  of  boric  acid  with  which  the  water  becomes  charged  is  suffi- 
cient to  constitute  this  a  profitable  source  of  supply.  To  obtain  the 
acid,  large  brick-work  basins  are  built  round  the  steam  jets  in  such 
a  manner  that  the  liquid  can  be  caused  to  flow  from  one  to  another. 
Water  is  placed  in  the  highest  basin,  and  after  the  steam  from  the 
fumaroles  beneath  it  has  blown  through  for  twenty-four  hours  the 
liquid  is  passed  on  to  the  second  basin,  and  a  fresh  supply  of  water 
is  run  into  the  first.  In  this  way  the  water  passes  on  through  a 
series  of  four  or  five  such  basins,  receiving  the  steam  of  the  soffioni 
for  twenty-four  hours  in  each.  The  muddy  liquor,  after  passing 
through  a  settling  reservoir,  is  concentrated  by  evaporation,  the 
heat  from  the  natural  steam  being  utilised.  The  concentrated 
liquor,  having  a  specific  gravity  about  1.07,  is  allowed  to  cool 
in  lead-lined  tanks  ;  and  the  crystals,  after  being  drained,  are 
dried  upon  the  floor  of  a  chamber,  also  heated  by  the  natural 
steam.  The  crude  boric  acid  thus  obtained  is  purified  by  recrys- 
tallisation. 

2-Q 


6iO  Inorganic  Chemistry 

Boric  acid  may  be  prepared  by  the  action  of  sulphuric  acid  or 
hydrochloric  acid  upon  a  strong  solution  of  borax  — 


Properties.  —  Boric  acid  crystallises  in  lustrous  white  laminae, 
which  are  soft  and  soapy  to  the  touch.  100  parts  of  water  at  18° 
dissolve  3.9  parts  of  the  acid.  The  aqueous  solution  turns  blue 
litmus  to  a  port  wine  red,  similar  to  the  colour  produced  by  car- 
bonic acid.  In  contact  with  turmeric  paper  it  gives  a  brown 
stain  resembling  that  caused  by  alkalies,  but  readily  distinguished 
by  not  being  destroyed  by  acids  and  by  being  turned  black  in 
contact  with  a  solution  of  sodium  hydroxide.  Boric  acid  is  more 
soluble  in  alcohol  than  in  water,  and  when  this  solution  is  boiled 
a  portion  of  the  boric  acid  volatilises  with  the  alcohol  and  imparts 
a  green  colour  to  the  flame  of  the  burning  vapour. 

The  orthoborates  are  mostly  unstable  salts. 

Metaborie  Aeid,  H2B2O4,  is  obtained  when  boric  acid  is  heated 
to  1  00°  — 

2H3BO3  =  2H2O  +  H2B2O4. 

The  metaborates  are  more  stable  salts  than  the  orthoborates. 
The  acid  is  dibasic,  and  forms  normal  and  acid  salts  as  well  as 
super-acid  salts,  thus  — 

Normal  potassium  metaborate     .         .     K2B2O4. 
Acid  potassium  metaborate          .         .     HKB2O4. 
Super-acid  potassium  metaborate         .     HKB2O4,H2B2O4. 

Pyroborie  Aeid,  H2B4O7,  is  obtained  by  heating  either  meta- 
boric  acid  or  orthoboric  acid  to  140°  for  some  time  — 

2H2B2O4=  H2O  +  H2B4O7. 
4H3BO3  =  5H2O  +  H2B4O7. 

Borax.  —  The  most  important  salt  of  pyroboric  acid  is  the  sodium 
salt,  ordinary  borax,  Na2B4O7,10H2O.  This  compound  occurs 
naturally  as  the  mineral  tincal.  It  is  manufactured  from  boric 
acid  by  double  decomposition  with  sodium  carbonate  — 


Anhydrous  sodium  carbonate  is  added  to  a  boiling  solution  of 


Boron   Trifluoride  611 

boric  acid,  and  the  liquid  is  then  allowed  to  crystallise,  when  it 
forms  large  transparent  prisms  belonging  to  the  mono-symmetric 
system  of  the  composition  Na2B4O7,10H2O. 

The  chief  source  of  borax,  however,  is  furnished  by  the  natural 
deposits  of  borate  of  lime  in  Bolivia.  The  powdered  mineral  is 
boiled  with  water,  and  soda  ash  is  added  to  the  mixture,  when 
calcium  carbonate  is  precipitated,  and  a  mixture  of  borax  and 
sodium  metaborate  is  formed — 

Ca2B6On  +  2Na2CO3  =  2CaCO3  +  Na2B4Or  +  Na2B2O4. 

On  crystallisation  the  borax  deposits,  and  the  more  soluble 
metaborate  remains  in  the  mother-liquor.  On  concentrating 
these  mother-liquors  and  blowing  carbon  dioxide  through  the 
solution,  the  metaborate  is  converted  into  borax,  which  is  pre- 
cipitated as  a  fine  meal,  leaving  sodium  carbonate  in  solution — 

2Na2B2O4  +  CO2  =  Na2CO3  +  Na2B4O7. 

When  heated,  borax  loses  its  water  of  crystallisation  and  swells 
up,  forming  a  white  porous  mass,  which  finally  melts  to  a  clear  glass. 

loo  parts  of  water  at  10°  dissolve  4.6  parts  of  crystallised 
borax,  and  at  100°,  201.4  parts;  the  solution  possesses  a  feeble 
alkaline  reaction. 

When  deposited  slowly  from  warm  solutions  (i.e.  above  about 
50°  C),  borax  crystallises  in  octahedrons  having  the  composition 
Na2B4OT,5H2O  ;  but  when  crystallised  without  any  special  precau- 
tions it  forms  prismatic  crystals  containing  10  molecules  of  water. 
This  is  the  ordinary  form  in  which  borax  is  met  with. 

Boron  Trifluoride,  BF3,  is  formed  when  boron  is  brought  into 
fluorine  ;  the  boron  takes  fire  spontaneously  in  the  gas,  forming 
the  trifluoride. 

It  is  also  produced  when  a  mixture  of  dry  powdered  fluorspar 
and  boron  trioxide  is  heated  to  redness  in  an  iron  vessel,  calcium 
borate  being  at  the  same  time  produced — 

2B2O3  +  3CaF2  =  Ca3B2O6  +  2B  F3. 

It  is  more  conveniently  prepared  by  heating  together  fluorspar, 
boron  trioxide,  and  sulphuric  acid.  The  reaction  may  be  regarded 
as  taking  place  in  two  stages,  thus — 

(i.) 
(2.) 


6i2  Inorganic  Chemistry 

Properties.  —  Boron  trifluoride  is  a  colourless,  pungent-smelling 
gas,  which  fumes  strongly  in  moist  air  on  account  of  its  powerful 
affinity  for  water.  So  great  is  this  affinity,  that  a  strip  of  paper 
introduced  into  the  gas  is  charred,  by  the  abstraction  of  the 
elements  of  water. 

Boron  fluoride  neither  burns  nor  supports  the  combustion  of 
ordinary  combustibles.  When  potassium  is  heated  in  the  gas  it 
burns  brilliantly,  forming  the  borofluoride. 

At  o°  one  volume  of  water  dissolves  about  1000  volumes  of  the 
gas,  the  absorption  being  attended  with  rise  of  temperature. 

When  the  gas  is  passed  into  water  until  the  solution  is  distinctly  acid,  a 
mixture  of  metaboric  acid  and  hydrofluoboric  acid  is  obtained;  the  former 
separates  out,  while  the  latter  remains  in  solution  — 

8BF3+4H20  =  H2B.204+6HBF4. 

When  the  gas  is  passed  into  water  until  the  latter  is  saturated,  a  syrup-like 
liquid  is  obtained  which  chars  organic  matter  and  is  strongly  corrosive.  This 
liquid  is  sometimes  called  fluoboric  acid,  and  contains  boron  trifluoride  and 
water  in  the  proportions  represented  by  the  formula  2BF3,4H2O  ;  or  it  may 
be  regarded  as  consisting  of  metaboric  acid  and  hydrofluoric  acid,  as  ex- 
pressed by  the  formula  H2B2O4,6HF.*  In  presence  of  an  excess  of  water, 
this  substance  is  decomposed  into  metaboric  acid  and  hydrofluoboric  acid. 

When  mixed  with  its  own  volume  of  dry  ammonia  gas,  boron  fluoride  forms 
a  white  crystalline  compound,  having  the  composition  represented  by  the 
formula  BF3,NH3.  This  substance  may  be  sublimed  without  change.  Two 
other  compounds  with  ammonia  are  known,  namely  BF3,2NH3,  and 
BF3,3NH3.  These  are  both  colourless  liquids,  which  on  being  heated  give  off 
ammonia,  leaving  the  solid  BF3,NH3. 

The  salts  of  hydrofluoboric  acid,  HBF4,  are  known  as  borofluorides  (sometimes 
ftuoborates],  and  are  formed  by  the  action  of  the  acid  upon  metallic  hydroxides  — 

HBF4+KHO  =  H20  +  KBF4. 

In  many  instances  their  aqueous  solutions  redden  litmus  ;  this  is  the  case 
with  ammonium  borofluoride,  NH4BF4,  and  calcium  borofluoride,  Ca(BF4)2. 

Boron  Trichloride,  BC13,  is  produced  when  boron  is  heated  in 
a  stream  of  dry  chlorine. 

It  is  most  readily  prepared  by  passing  dry  chlorine  over  an 
intimate  mixture  of  boron  trioxide  and  charcoal,  heated  to  redness 
in  a  porcelain  tube.  The  volatile  product  is  condensed  in  a  tube 
immersed  in  a  freezing-  mixture  — 


*  It  is  considered  very  doubtful  whether  this  substance  can  be  regarded 
as  a  definite  compound. 


Boron  Sulphide  613 

Properties.  —  Boron  trichloride  is  a  mobile,  colourless  liquid, 
boiling  at  18.23.  It  fumes  in  moist  air,  being  decomposed  in 
contact  with  water,  with  formation  of  boric  and  hydrochloric 
acids  — 

2O  =  B(HO)3 


Boron  trichloride  unites  directly  with  dry  gaseous  ammonia, 
with  evolution  of  considerable  heat,  forming  a  white  crystalline 
compound,  having  the  composition  2BC13,3NH3. 

Boron  Hydride,  BH3.  —  This  compound  has  never  been  obtained  in  a  state 
of  purity.  When  magnesium  boride  (an  impure  substance  obtained  by  fusing 
boron  trioxide  and  magnesium  in  a  covered  crucible)  is  acted  upon  by 
hydrochloric  acid,  a  gas  is  evolved  which  has  a  characteristic  and  unpleasant 
smell,  and  which  produces  headache  and  sickness  when  inhaled.  The  gas 
is  largely  hydrogen,  containing,  however,  a  certain  quantity  of  boron  hydride, 
which  imparts  to  the  flame  a  green  colour,  and  produces  boron  trioxide. 
When  passed  through  a  heated  tube,  boron  is  deposited  as  a  brown  film. 
When  burnt  with  a  limited  supply  of  air,  or  when  a  cold  porcelain  dish  is 
depressed  into  the  flame  of  the  burning  gas,  a  brown  stain  of  boron  is 
deposited. 

Boron  Nitride,  BN,  is  formed  when  boron  is  strongly  heated  in  nitrogen 
or  in  ammonia.  It  is  best  obtained  by  heating,  in  a  covered  platinum 
crucible,  a  mixture  of  one  part  of  dehydrated  borax,  and  two  parts  of 
ammonium  chloride  — 

Na.2B4O7  +  2N  H4C1  =  2BN  +  B2O3  +  2NaCl  +  4H2O. 

Boron  nitride  is  a  white  amorphous  powder.  It  is  insoluble  in  water,  but 
is  slowly  acted  upon  by  boiling  caustic  alkalies,  with  evolution  of  ammonia  — 

BN  +  3KHO=  K3BO3+  NH3. 

Heated  in  a  current  of  steam  it  forms  boron  trioxide  and  ammonia  — 
2BN  +  3H20=B.203+2NH3. 

Boron  Sulphide,  B2S3,  is  prepared  by  heating  a  mixture  of  boron  trioxide 
and  carbon  (made  by  mixing  boron  trioxide  and  soot  with  oil,  and  heating 
the  pellets  out  of  contact  with  air)  to  bright  redness  in  a  stream  of  vapour 
of  carbon  disulphide  — 


Boron  sulphide  is  a  yellowish  solid,  consisting  of  small  crystals.  It  has 
a  strong  unpleasant  smell,  and  its  vapour  attacks  the  eyes.  It  is  immediately 
decomposed  by  water,  being  converted  into  boric  acid  and  sulphuretted 
bvdrogen  — 

B2S3+6H20=2B(HO)3+3H2S. 


6  14  Inorganic  Chemistry 

ALUMINIUM. 

Symbol,  Al.     Atomic  weight  =  27.1. 

Oeeurrenee.  —  Aluminium  is  one  of  the  most  abundant  of  all 
the  elements,  although  it  has  never  been  found  in  the  uncombined 
state.  In  combination  with  oxygen  as  A12O3,  it  constitutes  such 
minerals  as  corundum,  ruby,  sapphire.  As  the  hydrated  oxide, 
A12O3,2H2O,  it  occurs  associated  with  iron  oxide  in  the  mineral 
bauxite,  which  constitutes  the  chief  source  from  which  the  metal 
itself  is  obtained.  As  a  double  fluoride  of  aluminium  and  sodium, 
AlF3,3NaF,  it  occurs  in  the  mineral  cryolite,  and  as  a  hydrated 
phosphate  in  the  various  forms  of  turquoise.  Aluminium  is  met 
with  in  enormous  quantities  in  the  form  of  silicate,  constituting 
the  various  clays  ;  and  as  compound  silicates  in  the  felspars,  and 
other  common  minerals  constituting  a  large  proportion  of  the 
solid  crust  of  the  earth. 

Mode  of  Formation.—  Prior  to  the  advent  of  the  electric 
furnace  as  a  manufacturing  agent,  aluminium  was  obtained  from 
the  mineral  bauxite  by  the  following  method  :  —  The  process  was 
conducted  in  four  stages  —  (i.)  and  (2.)  The  preparation  of  pure 
aluminium  oxide,  free  from  iron.  (3.)  The  preparation  of  a  double 
chloride  of  aluminium  and  sodium.  (4.)  The  reduction  of  the 
double  chloride  by  means  of  sodium. 

(i.)  The  powdered  bauxite  (usually  containing  about  50  per 
cent,  of  alumina)  was  mixed  with  sodium  carbonate  and  heated  for 
five  or  six  hours  in  a  reverberatory  furnace,  when  carbon  dioxide 
is  evolved  and  sodium  aluminate  is  formed  — 

A12O3  +  Na2CO3  =  2NaAlO2+  CO2. 

(2.)  The  sodium  aluminate  was  extracted  with  water,  leaving  the 
iron  in  the  form  of  insoluble  oxide.  Through  the  filtered  liquid 
a  stream  of  carbon  dioxide  was  then  passed,  which  decomposes 
sodium  aluminate,  regenerating  sodium  carbonate,  and  precipi- 
tating hydrated  aluminium  oxide  — 


(3.)  The  purified  alumina,  after  being  washed  and  dried,  was  mixed 
with  sodium  chloride  and  powdered  wood  charcoal,  and  sufficient 
water  added  to  enable  the  mixture  to  be  worked  up  into  balls, 


Aluminium 


These  were  dried  and  packed  into  a  vertical  fireclay  cylinder,  and 
strongly  heated  in  a  stream  of  chlorine  — 


2  =  3CO  +  2A1C13. 


The  aluminium  chloride  combines  with  the  sodium  chloride 
present  in  the  mixture,  forming  the  double  chloride,  AlCl3,NaCl, 
which  volatilises  from  the  retort,  and  is  condensed  in  a  receiver. 


FIG.  146. 

(4.)  Finally  the  double  chloride  of  aluminium  and  sodium  was 
strongly  heated  with  metallic  sodium  and  powdered  cryolite  (to 
serve  as  a  flux) — 

AlCl3,NaCl  +  3Na  =  4NaCl  +  Al. 

Electrolytic  Method. — At  the  present  time  aluminium  is  ex- 
clusively obtained  by  means  of  the  electric  furnace.  The  process 
is  an  electrolytic  one,  the  electrolyte  being  a  solution  of  alumina 
in  a  bath  of  molten  cryolite.  One  of  the  most  modern  forms  of 


616  Inorganic  Chemistry 

apparatus  for  the  purpose  (Borcher's)  is  shown  in  section  in 
Fig.  146. 

It  consists  of  an  iron  cylinder  or  crucible  C,  with  a  fireclay 
bottom  F,  and  thickly  lined  throughout  with  alumina,  L.  The 
cathode  consists  of  a  steel  plate  S,  let  into  the  bottom  of  the 
crucible,  into  which  is  screwed  the  copper  tube  T.  To  prevent 
the  steel  plate  from  becoming  too  much  heated,  and  in  consequence 
combining  with  the  aluminium,  an  arrangement  is  made  to  cir- 
culate water  through  the  pipe  T. 

The  anode  consists  of  a  thick  carbon  rod,  or  bundle  of  rods, 
which  can  be  raised  or  lowered  at  will.  A  few  fragments  of 
aluminium,  together  with  a  small  quantity  of  cryolite,  is  first  placed 
in  the  crucible,  and  melted  by  bringing  the  anode  down  upon  it. 
The  fused  button  of  aluminium  then  becomes  the  cathode.  The 
crucible  is  then  gradually  filled  up  with  its  charge  of  cryolite  and 
bauxite  until  the  entire  mass  is  in  a  molten  state.  The  aluminium 
oxide  alone  is  decomposed  in  the  process,  the  oxygen  escaping 
through  an  opening  in  the  lid,  while  the  metal  collects  at  the  bottom 
and  is  drawn  off  at  the  tap-hole.  Fresh  bauxite  is  added  in  small 
quantities  at  a  time  as  the  action  continues.  It  is  found  that  the 
lining  of  the  crucible,  although  of  alumina,  is  not  dissolved  to  a 
very  great  extent,  owing  to  the  cooling  of  the  surface  by  outside 
exposure  to  the  air. 

Properties. — Aluminium  is  a  tin-white  metal,  possessing  great 
tensile  strength.  It  is  very  ductile  and  malleable,  but  requires 
frequent  annealing  during  the  process  of  drawing  or  hammering. 
Its  specific  gravity  is  2.58  ;  by  hammering  and  rolling  it  may  be 
raised  to  2.68.  Its  power  of  conducting  heat  and  electricity  is 
about  one-third  that  of  silver.  Aluminium  is  an  extremely  sono- 
rous metal,  and  when  struck  it  emits  a  clear  and  sustained  note. 
It  is  not  tarnished  by  air  under  ordinary  circumstances,  but  when 
strongly  heated  it  becomes  oxidised  ;  and  in  the  condition  of  thin 
foil  it  readily  burns  in  oxygen,  forming  alumina,  A12O3.  The  metal 
melts  at  a  temperature  about  655°.  Aluminium  is  scarcely  acted 
upon  by  nitric  acid  of  any  strength,  but  readily  dissolves  in  hydro- 
chloric acid,  and  in  solutions  of  sodium  or  potassium  hydroxide 
with  elimination  of  hydrogen.  When  heated  with  strong  sulphuric 
acid,  aluminium  sulphate  is  formed,  and  sulphur  dioxide  is 
evolved. 

Organic  acids  are  almost  without  action  upon  aluminium,  but 
in  the  presence  of  sodium  chloride  they  are  capable  of  dissolv- 


Aluminium  Oxide  617 

ing  it  to  a  slight  extent.  Pure  aluminium  is  scarcely  acted  upon 
by  water  or  steam,  but  the  presence  of  impurities  such  as  usually 
occur  in  the  commercial  metal  renders  it  much  more  readily 
oxidised. 

Aluminium  is  a  highly  electro-positive  element,  and  is  capable 
of  reducing  a  number  of  other  metals  from  their  combinations  with 
oxygen  or  sulphur.  Thus,  when  finely  divided  aluminium  is  heated 
with  the  oxides  of  such  metals  as  manganese,  chromium,  tungsten, 
uranium,  along  with  lime  to  form  a  slag,  an  energetic  action  takes 
place,  in  which  the  aluminium  combines  with  the  oxygen,  and  the 
metals  are  thrown  out  of  combination,  and  are  obtained  as  a 
coherent  mass.  Similarly,  iron  pyrites  is  reduced  to  the  condition 
of  metallic  iron,  with  the  formation  of  aluminium  sulphide. 

The  extreme  readiness  with  which  aluminium  is  able  to  effect  such 
reduction,  and  the  exceedingly  high  temperature  which  is  reached 
by  the  action,  have  led  to  some  useful  applications,  such  as  the 
welding  of  iron,  &c.  The  property  may  readily  be  demonstrated 
by  heating  upon  a  spatula  or  an  iron  plate  a  small  quantity  of  a 
mixture  of  powdered  aluminium  and  copper  oxide.  When  such  a 
mixture  is  brought  into  a  Bunsen  flame  the  action  takes  place 
immediately,  and  is  accompanied  by  an  instantaneous  and  vivid 
flash  of  light. 

Alloys  of  Aluminium. — The  most  important  of  these  is  an 
alloy  with  copper,  containing  10  per  cent,  of  aluminium,  and 
known  as  aluminium  bronze.  This  alloy  has  a  yellow  colour, 
resembling  that  of  gold  ;  it  is  scarcely  tarnished  by  exposure  to 
air,  and  is  susceptible  of  a  high  polish.  Its  specific  gravity  is  7.69, 
and  it  possesses  a  tenacity  equal  to  that  of  steel,  and  more  than 
twice  that  of  the  best  gun-metal.  The  alloy  is  malleable,  and 
yields  good  castings,  and  on  account  of  its  many  valuable  pro- 
perties it  is  employed  for  a  variety  of  purposes. 

Aluminium  Oxide  (alumina),  A12O3,  occurs  native  in  a  colour- 
less crystalline  condition  as  corundum,  and  coloured  by  traces  of 
various  metallic  oxides  in  such  precious  stones  as  ruby,  sapphire, 
and  amethyst.  In  a  less  pure  condition,  it  occurs  in  large  quantities 
as  emery.  These  naturally  occurring  crystalline  forms  of  alumina 
are  extremely  hard,  ranking  second  only  to  diamond.  Alumina  is 
obtained  in  an  amorphous  condition  by  igniting  either  the  pre- 
cipitated hydroxide  or  ammonia  alum,  thus — 

2A1(HO)3  =  3 


618  Inorganic  Chemistry. 

It  is  also  obtained  by  the  action  of  carbon  dioxide  upon  sodium 
aluminate  (p.  614).  In  the  crystalline  form  it  is  obtained  by 
strongly  heating  a  mixture  of  aluminium  fluoride  and  boron  tri- 
oxide  — 

2A1F3  +  B2O3  =  A12O3  +  2B  F3. 

The  boron  trifluoride  volatilises,  leaving  alumina  in  the  form  of 
rhombohedral  crystals.  Artificial  rubies  have  been  obtained  by 
heating  barium  fluoride  with  alumina,  and  adding  a  trace  of 
potassium  dichromate. 

Amorphous  alumina  is  a  soft  white  powder,  insoluble  in  water, 
but  dissolved  by  acids  with  the  formation  of  aluminium  salts  ; 
after  being  strongly  heated,  however,  alumina  is  attacked  only 
with  slowness  by  hydrochloric  or  sulphuric  acid. 

Aluminium  Hydroxides.—  When  ammonia  is  added  to  a 
solution  of  an  aluminium  salt,  a  white  gelatinous  precipitate  is 
obtained,  consisting  of  the  trihydrate,  A12O33H2O,  or  A1(HO)3. 
If  this  be  heated  to  300°  it  loses  water,  and  is  converted  into  the 
mono-hydrate,  A12O3,H2O,  or  AIO(HO).  If  the  precipitation  be 
made  in  boiling  solution,  or  if  the  trihydrate  be  heated  to  100°, 
a  compound  is  obtained  having  a  composition  expressed  by  the 
formula  A12O3,2H2O,  or  A12O(HO)4. 

These  compounds  are  soluble  in  acids,  and  all  yield  the  same 
aluminium  salts. 

Aluminium  hydroxide  unites  with  many  soluble  organic  colour- 
ing-matters, and  precipitates  them  from  solution  as  lakes.  Upon 
this  property  depends  the  use  of  aluminium  salts  as  mordants  in 
dyeing  and  calico-printing  :  the  colouring-matter  being  held  in  the 
fibres  of  the  material  by  the  aluminium  hydroxide,  which  is  pre- 
viously precipitated  upon  the  fabric. 

Aluminates.  —  Alumina  is  capable  of  acting  as  a  feeble  acidic 
oxide  :  thus,  the  hydroxides  are  dissolved  by  sodium  or  potassium 
hydroxide,  yielding  salts  known  as  aluminates.  Certain  alu- 
minates  occur  native,  such  as  spinelle  (magnesium  aluminate), 
Al2O3,MgO,  or  Mg(AlO2)2,  and  chrysoberyl  (beryllium  aluminate), 
Al2O3,BeO,  or  Be(AlO2)2.  Sodium  aluminate  is  manufactured  by 
fusing  bauxite  with  sodium  carbonate  (p.  614),  or  by  boiling 
powdered  cryolite  with  milk  of  lime— 


It  is  also  produced  when  a  mixture  of  cryolite  and  lime  is  heated 


Aluminium  Sulphate  619 

to  redness.  Under  these  conditions  the  0r///0aluminate  is  formed, 
which  when  acted  upon  by  water  undergoes  hydrolysis  into  the 
w^/aluminate  and  free  alkali  — 

AlF3,3NaF  +  3CaO  =  3CaF9  +  Na3AlO3. 
Na3A103+  H,O  -  NaAlO2  +  2NaHO. 

Sodium  aluminate  is  largely  employed  as  a  mordant  in  dyeing 
and  calico-printing  owing  to  the  ease  with  which  it  is  decomposed, 
yielding  alumina.  Thus  it  is  hydrolysed  to  a  considerable  extent 
by  water,  the  action  being  greatly  accelerated  by  the  addition  of 
a  small  quantity  of  aluminium  hydroxide  — 


It  is  decomposed  by  carbon  dioxide  (p.  614)  and  also  by  any 
halogen  compound  of  aluminium  :  powdered  cryolite  being  often 
employed  for  this  purpose  — 


Aluminium  Sulphate,  A12(SO4)3,18H2O,  is  found  native  as  the 
minerals  hair  salt  and  aluminite^  the   latter  being  a  basic   salt 


FIG.  147- 

having  the  composition  A12O3SO3,9H2O.  Large  quantities  of 
commercial  aluminium  sulphate  are  made  by  directly  dissolving 
bauxite  in  sulphuric  acid.  The  product,  however,  contains  iron, 


62O  Inorganic  Chemistry 

which  is  detrimental  to  the  technical  uses  to  which  the  sulphate  is 
applied,  and  from  which  therefore  it  must  be  carefully  purified. 
Pure  aluminium  sulphate  is  prepared  by  dissolving  the  hydroxide 
in  sulphuric  acid.  It  forms  a  white  difficultly  crystallisable  solid. 

The  Alums. — Aluminium  sulphate  unites  with  certain  other 
sulphates,  forming  double  salts,  which  belong  to  a  class  of  com- 
pounds known  as  the  alums.  The  most  important  of  these 
compounds  is  the  double  sulphate  of  aluminium  and  potassium, 
A12(SO4)3,K2SO4,24H2O,  known  as  potassium  alum,  or  simply 
alum. 

The  alums  have  the  general  formula  R2(SO4)3,M2SO4,24H2O, 
in  which  R  may  be  either  aluminium,  iron,  chromium,  manganese 
(indium  or  gallium),  and  M  a  monovalent  element  or  group,  such 
as  sodium,  potassium,  or  ammonium. 

These  compounds  are  all  isomorphous,  crystallising  in  the 
regular  system  (usually  in  cubes  or  octahedra)  with  twenty-four 
molecules  of  water.  Fig.  147  represents  a  crystal  of  potassium 
alum  (A)  and  potassium  chromium  alum  (B).  In  naming  the 
alums  *  it  is  usual,  when  the  salt  contains  aluminium,  only  to 
introduce  the  name  of  the  monovalent  element  or  group  :  thus, 
ammonium  alum,  or  potassium  alum,  signifies  the  double  sulphate 
of  ammonium,  or  potassium,  and  aluminium.  If,  on  the  other  hand, 
the  compound  contains  no  aluminium,  the  names  of  both  metals 
are  used,  thus,  potassium  chromium  alum,  ammonium  iron  alum. 

A  second  class  of  double  sulphates  is  known,  which  resemble  the  alums, 
although  they  are  not  isomorphous  with  them.  These  are  termed  pseudo- 
alums.  They  may  be  regarded  as  alums,  in  which  the  two  atoms  of  the 
monovalent  element  are  replaced  by  one  atom  of  a  divalent  element,  thus — 

Manganese  aluminium,  pseudo-alum      .         .  Al2(SO4)3MnSO4,24H2O. 

Iron  aluminium,  pseudo-alum         .         .         .  Al2(SO4)3FeSO4,24H2O. 

Copper  iron,  pseudo-alum      ....  Fe2(SO4)3CuSO4,24H2O. 

Zinc  iron,  pseudo-alum Fe2(SO4)3ZnSO4,24H2O. 

Magnesium  manganese,  pseudo-alum     .         .  Mn2(SO4)3MgSO4,24H2O. 

The  alums  are  all  soluble  in  water,  and  their  solutions  have  an 

*  Selenic  acid  (the  selenium  analogue  of  sulphuric  acid)  forms  a  similarly 
constituted  series  of  double  selenates,  crystallising  in  the  same  form,  and 
with  the  same  number  of  molecules  of  water.  The  system  of  nomenclature 
adopted  for  these  compounds  is  the  same :  thus,  ammonium  selenio-alum 
signifies  the  double  selenate  of  ammonium  and  aluminium,  while  potassium 
chromium  selenio-alum  represents  the  double  selenate  of  potassium  and 
chromium. 


Alum  621 

acid  reaction  and  possess  an  astringent  taste.  When  heated 
they  gradually  part  with  water,  and  at  higher  temperatures  are 
broken  up  into  oxides  and  alkaline  sulphates ;  in  the  case  of 
ammonium  alums,  leaving  only  the  metallic  oxide. 

Potassium  Alum,  A12(SO4)3,K2SO4,24H2O,  is  prepared  by 
the  addition  of  the  requisite  quantity  of  potassium  sulphate  to 
aluminium  sulphate.  A  considerable  quantity  of  alum  is  also 
obtained  from  a  naturally  occurring  basic  potassium  alum,  known 
as  alum  stone^  or  alunite,  which  has  the  composition  A12(SO4)3, 
K2SO4,2A12O3,8H2O.  At  Tolfa  this  is  first  calcined,  and  after- 
wards lixiviated  with  water,  which  dissolves  the  potassium  alum, 
leaving  alumina  undissolved.  The  alum  so  obtained  is  known  as 
Roman  alum;  and  although  it  has  a  reddish  colour,  due  to  the  pre- 
sence of  iron,  this  iron  is  present  only  as  the  insoluble  oxide,  which 
is  readily  removed,  and  the  salt  is  in  reality  extremely  pure. 

Alunite  is  also  converted  into  alum,  by  treating  the  calcined 
mineral  with  sulphuric  acid,  and  adding  the  requisite  quantity  of 
potassium  sulphate.  A  large  quantity  of  alum  is  manufactured 
from  alum  shale,  which  is  a  bituminous  mineral,  consisting  chiefly 
of  aluminium  silicate,  with  finely  divided  iron  pyrites  dissemi- 
nated throughout  the  mass.  The  shale  is  usually  first  roasted, 
and  is  then  exposed  to  the  action  of  air  and  moisture,  whereby 
the  oxidation  of  the  pyrites  is  completed.  The  result  of  this 
oxidation  is  the  formation  of  sulphuric  acid,  which,  acting  upon 
the  aluminium  silicate,  forms  aluminium  sulphate,  while  the  iron 
is  converted  into  ferrous  and  ferric  sulphates,  and  ferric  oxide.  The 
oxidised  mass  is  then  lixiviated  with  water,  and,  after  concentra- 
tion, the  requisite  quantity  of  potassium  chloride  or  sulphate  is 
added  to  the  hot  liquor.  (The  use  of  potassium  chloride  is  pre- 
ferable, as  by  double  decomposition  the  ferrous  and  ferric  sulphates 
are  converted  into  the  very  soluble  chlorides,  and  an  equivalent 
amount  of  potassium  sulphate  is  formed.)  The  liquor  is  stirred 
mechanically  during  its  cooling,  whereby  the  alum  is  deposited  in 
small  crystals  known  as  alum  meal,  which  permit  of  its  more 
ready  purification  by  recrystallisation. 

Alum. crystallises  in  fine  colourless  regular  octahedra,  which,  on 
exposure  to  the  air,  become  coated  with  a  white  efflorescence,  due 
not  to  loss  of  water,  but  to  absorption  of  atmospheric  ammonia,  and 
the  formation  of  a  basic  salt. 

The  solubility  of  alum  in  water  increases  rapidly  with  rise  of 
temperature.  Thus,  100  parts  of  water  at  o°  dissolve  3.9  parts  of 


622  Inorganic  Chemistry 

alum;  at  50°,  44.1  parts;  and  at  100°,  357.5  parts.  Alum  is  in- 
soluble in  alcohol. 

When  heated  to  42°,  alum  loses  1 1  molecules  of  water ;  and 
when  heated  to  61°  in  a  closed  vessel  over  sulphuric  acid,  it  parts 
with  1 8  molecules. 

On  the  application  of  heat,  alum  first  melts  in  its  own  water  of 
crystallisation,  which  is  gradually  expelled,  until  at  a  dull  red  heat 
the  salt  is  converted  into  a  white  porous  mass,  known  as  burnt 
alum.  At  a  still  higher  temperature  it  is  broken  up  into  potassium 
sulphate,  alumina,  and  sulphur  trioxide.  Burnt  alum  is  only  very 
slowly  dissolved  by  water.  The  chief  use  of  alum  is  as  a  mordant 
in  dyeing,  alum  being  a  salt  which  is  much  more  easily  obtained 
in  a  state  of  purity  than  aluminium  sulphate.  By  the  addition  of 
sodium  hydroxide  or  carbonate  to  a  solution  of  alum,  until  the 
precipitate  first  thrown  down  is  just  redissolved,  a  basic  alum  is 
produced  known  as  neutral  alum — 

2Al2(S04)3,K2S04+6NaHO  =  A]2(S04)3,Al2(HO)6,K2SO,+ 
3Na2SO4-}-K2SO4. 

This  solution  gives  up  its  alumina  to  the  fabric  with  great  ease, 
and  on  this  account  is  used  by  dyers  and  calico-printers  as  a 
mordant. 

When  this  solution  is  heated  to  40°,  ordinary  alum  is  reformed, 
and  a  precipitate  is  obtained  consisting  of  another  basic  salt,  hav- 
ing the  same  composition  as  alunite^  thus — • 

2A12(S04)3,A12(HO)6,K2S04=A12(S04)3,K2S04)  + 
A12(S04)3,2A1203,K2S04  +  6H20. 

Aluminium  Fluoride,  A1F3.— This  compound  may  be  prepared  by  passing 
gaseous  hydrochloric  acid  over  a  mixture  of  fluorspar  and  alumina  heated  to 
whiteness  in  a  graphite  tube,  when  aluminium  fluoride  volatilises,  leaving 
calcium  chloride — 

3CaF2  +  A12O3  +  6HC1 = 3H2O  +  3CaCl2  +  2A1F3. 

In  the  form  of  a  crystalline  hydrate  it  may  be  obtained  by  dissolving  alumina 
in  aqueous  hydrofluoric  acid — 

A1203  +  6HF  +  H20=2A1F3,7H20. 

Aluminium  fluoride  forms  colourless  rhombohedral  crystals,  which  are  in- 
soluble in  water.  It  combines  with  alkali  fluorides,  forming  insoluble  double 
fluorides,  of  which  the  sodium  compound  is  the  most  important,  AlF3,3NaF. 
This  compound  occurs  native  as  the  mineral  cryolite. 


Thallium  623 

Aluminium  Chloride,  A1C13.  —  This  compound  is  produced 
when  powdered  aluminium  is  strongly  heated  in  chlorine,  or  with 
certain  metallic  chlorides,  such  as  zinc  chloride.  It  is  best 
obtained  by  passing  chlorine  over  a  strongly  heated  mixture  of 
alumina  and  charcoal. 

An  aqueous  solution  of  aluminium  chloride  may  be  obtained  by 
dissolving  alumina  in  hydrochloric  acid.  On  evaporation  the 
solution  deposits  crystals  of  a  hydrate,  A1C13,6H2O. 

Aluminium  chloride  forms  white  hexagonal  crystals,  which  fume 
strongly  in  moist  air.  When  gently  heated  it  vaporises,  and  sub- 
limes without  fusion.  When  heated  under  pressure  of  its  own 
vapour,  the  compound  melts.  It  dissolves  in  water  with  the 
evolution  of  heat,  and  the  solution  on  evaporation  deposits  the 
hydrated  chloride,  which,  on  being  heated,  breaks  up  into  hydro- 
chloric acid,  water,  and  alumina  — 


A1C13,6H2O  =  3HC1  +  3H20  +  A1(HO)3. 

Aluminium  chloride  unites  with  other  metallic  chlorides,  forming 
double  salts,  of  which  the  sodium  compound  AlCl3,NaCl  (p.  615) 
is  the  most  important.  It  also  combines  with  ammonia,  forming 
the  compounds  A1C13,6NH3  and  A1C13,NH3. 

Aluminium  Sulphide,  A12S3.  —  When  finely  divided  aluminium 
is  heated  with  iron  pyrites,  an  energetic  reaction  takes  place  ; 
metallic  iron  being  reduced,  and  aluminium  sulphide  being  formed. 
The  same  compound  is  produced  when  sulphur  is  thrown  upon 
strongly  heated  aluminium.  As  obtained  by  these  methods, 
aluminium  sulphide  is  a  greyish-black  solid,  which,  when  thrown 
into  water,  is  converted  into  the  oxide  with  evolution  of  sul- 
phuretted hydrogen  — 


The  compound  is  decomposed  in  the  same  manner  by  atmos- 
pheric moisture  when  exposed  to  the  air. 


THALLIUM. 

Formula,  Tl.     Atomic  weight  =  204. 

History. — Thallium  was  discovered  by  Crookes  (1861)  in  the 
seleniferous  deposit  from  a  sulphuric  acid  manufactory.     In  the 


624  Inorganic  Chemistry 

spectroscopic  examination  of  certain  residues  obtained  in  the  ex- 
traction of  selenium  from  this  deposit,  the  presence  of  an  unknown 
element  was  manifested,  by  the  appearance  of  one  bright  green 
line.  From  its  characteristic  spectrum,  the  name  thallium  (signi- 
fying a  green  twig)  was  given  to  the  element. 

Occurrence. — Thallium  is  found  in  small  quantities  in  many 
varieties  of  iron  pyrites,  and  when  these  are  employed  in  the 
manufacture  of  sulphuric  acid,  oxide  of  thallium  collects  in  the 
flue  dust  of  the  pyrites  burners.  Thallium  also  occurs  associated 
with  copper,  selenium,  and  silver,  in  the  rare  mineral  crookesite. 

Mode  Of  Formation. — The  metal  is  obtained  by  reducing  the 

sulphate,   by   immersing   strips   of  zinc   into  the  solution.      The 

,  thallium  is  deposited  upon  the  zinc,  as  a  spongy  or  crystalline  mass, 

which   is   then   pressed   together  and   fused   beneath   potassium 

cyanide  in  a  crucible. 

Properties. — Thallium  is  a  soft  heavy  metal,  resembling  lead. 
It  is  readily  cut  with  a  knife,  and  leaves  a  streak  when  drawn 
across  paper.  When  preserved  out  of  contact  with  air  it  is  a  tin- 
white  lustrous  metal  ;  but  on  exposure  to  the  air  it  tarnishes 
upon  its  surface,  with  the  formation  of  black  thallous  oxide.  Its 
specific  gravity  is  n.8,  and  it  melts  at  290°. 

When  exposed  to  air  and  moisture,  or  when  placed  in  water 
which  is  free  to  absorb  atmospheric  oxygen,  the  metal  is  slowly 
converted  into  thallous  hydroxide,  which  is  soluble  in  water,  and 
imparts  to  the  liquid  a  strong  alkaline  reaction.  The  solution 
absorbs  carbon  dioxide,  with  the  formation  of  thallous  carbonate. 
When  heated  in  the  air  thallium  melts,  and  rapidly  oxidises  to 
thallium  trioxide,  T12O3  ;  heated  in  oxygen  it  burns,  forming  the 
same  oxide.  It  readily  burns  when  heated  in  chlorine,  producing 
thallous  chloride,  T1C1.  The  metal  is  soluble  in  dilute  acids. 

Oxides  Of  Thallium.— Two  oxides  are  known,  namely,  thallous 
oxide,  T12O,  and  thallic  oxide,  T12O3. 

Thallous  Oxide,  T12O,  forms  as  a  dark  grey  film  upon  the 
surface  of  the  metal,  on  exposure  to  the  air.  It  may  also  be 
obtained  by  heating  the  hydroxide  to  100°.  It  dissolves  in  water, 
forming  the  hydroxide. 

Thallous  Hydroxide  is  obtained  by  the  addition  of  barium 
hydroxide  to  a  solution  of  thallous  sulphate,  the  precipitated  barium 
sulphate  being  removed  by  filtration — 

Tl2SO4  +  Ba(HO)2  = 


Thallic  Chloride  625 

The  solution,  on  concentration,  deposits  yellow  needle-shaped 
crystals  of  T1HO,H2O.  Thallous  hydroxide  is  soluble  in  water, 
yielding  an  alkaline  solution  which  gives  a  brown  stain  upon 
turmeric  paper.  The  stain  soon  disappears,  owing  to  the  de- 
struction of  the  colouring-matter,  and  is  thereby  distinguished 
from  the  similar  stains  produced  by  sodium  and  potassium 
hydroxides. 

Thallie  Oxide,  T12O3,  is  obtained  when  thallium  burns  in  the 
air,  or  when  thallium  oxyhydroxide,  TIO(HO),  is  heated  to  100°. 
It  forms  a  dark  reddish  powder,  insoluble  in  water.  In  warm 
dilute  sulphuric  acid  it  dissolves,  forming  thallic  sulphate  — 

T12O3  +  3H2SO4  =  T12(SO4)3  +  3H2O, 

but  with  hot  concentrated  acid  oxygen  is  evolved,  and  thallous 
sulphate  formed  — 

T12O3  +  H2SO4  =  T12SO44-O2  +  H2O. 

At  a  red  heat  thallic  oxide  is  converted  into  thallous  oxide  with 
loss  of  oxygen. 

Thallium  Oxyhydroxide,  TIO(HO),  is  formed  by  the  action  of 
potassium  hydroxide  upon  thallium  trichloride  — 

T1C13  +  3KHO  =  3KC1  +  H2O  +  T1O(HO). 

Thallous  Chloride,  T1C1,  is  obtained  as  a  white  curdy  precipi- 
tate when  hydrochloric  acid  is  added  to  a  solution  of  a  thallous 
salt.  It  is  considerably  more  soluble  in  hot  than  in  cold  water  : 
ico  parts  of  water  at  16°  dissolve  0.265  Part  \  and  at  100°,  1.427 
part  of  thallous  chloride. 

Thallic  Chloride,  T1C13,  is  formed  by  passing  chlorine  through 
water  in  which  thallous  chloride  is  suspended.  The  solution  so 
obtained,  on  evaporation  in  vacuo,  deposits  colourless  transparent 
crystals  of  T1C13,2H2O. 

When  either  thallium  or  thallous  chloride  is  gently  heated  in  a 
stream  of  chlorine,  a  compound  is  obtained,  having  the  composition 
T1C13,T1C1,  or  T12C14.  If  this  be  further  heated,  it  loses  chlorine, 
and  is  converted  into  a  yellow  crystalline  compound  of  the  com- 
position T1C13,3T1C1,  or  T14C16,  thus— 


2T12C14=C12 

2  R 


626  Inorganic  Chemistry 

ThallOUS  Oxy salts. —The  sulphate  T12SO4,  and  nitrate  T1NO3, 
are  best  obtained  by  dissolving  the  metal  in  the  respective  acids. 
Both  salts  are  soluble  in  water. 

ThallOUS  Carbonate,  T12CO3,  is  prepared  by  saturating  a  solu- 
tion of  thallous  hydroxide  with  carbon  dioxide.  The  salt  forms 
long  white  prismatic  (monosymmetric)  crystals,  which  are  mode- 
rately soluble  in  water,  giving  an  alkaline  solution. 

ThallOUS  Phosphate,  T13PO4,  is  obtained  by  precipitation  from 
a  thallous  solution,  by  the  corresponding  potassium  phosphate. 
The  monohydrogcn  phosphate,  HT12PO4,  on  being  heated  to  200°, 
is  converted  into  pyrophosphate — 

2HT12PO4= H2O  +  T14P2O7, 

and  the  dihydrogen  salt,  on  being  ignited,  yields  the  metaphos- 
phate — 

H2T1PO4=H2O  +  T1PO3. 

Thallie  OxysaltS.— The  chief  of  these  are  thallic  sulphate, 
T12(SO4)3 ;  and  thallic  nitrate,  T1(NO3)3.  They  are  obtained  by  the 
action  of  sulphuric  acid  and  nitric  acid  respectively  upon  thallic 
oxide  T12O3.  Thallic  sulphate  forms  colourless  crystals  of  the 
composition  T12(SO4)3,7H2O.  It  is  decomposed  by  excess  of  water, 
with  precipitation  of  the  hydrated  oxide  ;  and  when  heated  yields 
thallous  sulphate,  sulphur  trioxide,  and  oxygen — 

T12(SO4)3=T12SO4  +  2SO3  +  O,. 

Thallic  nitrate  is  depositedin  colourless  crystalsof  T1(NO3)3,8H2O, 
which  are  decomposed  in  the  presence  of  much  water. 


CHAPTER  IX 
THE  ELEMENTS  OF  GROUP  IV 


Family  A. 

Family  B. 

Titanium,  Ti     . 

.       48.1 

Carbon,  C    . 

12.00 

Zirconium,  Zr    . 

90.6 

Silicon,  Si    . 

28.3 

Cerium,  Ce 

.     140.25 

Germanium,  Ge  . 

72-5 

Thorium,  Th      . 

•     232.5 

Tin,  Sn 

119 

Lead,  Pb     . 

207.1 

Family  A  consists  of  four  rare  elements.*  Titanium,  as  the 
oxide  TiO2,  occurs  in  the  three  rare  minerals — rutile,  brookite,  and 
anastase.  The  metal  is  extremely  difficult  to  isolate  in  a  pure 
state,  owing  to  the  fact  that  it  unites  directly  with  nitrogen,  forming 
a  nitride. 

Zirconium  is  met  with  as  the  silicate  ZrSiO4  (or  ZrO9,SiO2)  in 
the  mineral  zircon.  Like  silicon,  it  has  been  obtained  in  two 
forms,  crystalline  and  amorphous.  The  latter  variety,  when  gently 
heated,  burns  in  the  air,  while  the  crystalline  variety  requires  the 
high  temperature  of  the  oxyhydrogen  flame  for  its  ignition. 

Cerium  occurs  associated  with  lanthanum  in  the  rare  minerals 
cerite  and  orthite,  and  with  yttrium  and  ytterbium  in  gadolinite  and 
wohlerite. 

Thorium  is  found  in  the  extremely  rare  minerals,  thorite  and 
orangeite,  met  with  in  Norway. 

Family  B. — In  this  family  the  rare  element  germanium  forms 
a  link  between  carbon  and  silicon  on  the  one  hand,  and  tin  and 
lead  on  the  other. 

Carbon  (the  typical  element)  is  essentially  non-metallic,  and 
forms  an  acidic  oxide.  Silicon  approaches  more  nearly  to  the 
metals  in  its  physical  properties,  but  its  oxide  is  still  acidic,  and 
but  few  compounds  are  known  in  which  silicon  functions  as  a  basic 
element.  Germanium  is  both  metallic  and  non-metallic  ;  its  oxide 

*  For  descriptions  of  these  rare  elements  the  student  is  referred  to  larger 
treatises. 

627 


628  Inorganic  Chemistry 

unites  with  acids  ;  and  it  also  combines  with  alkaline  hydroxides, 
forming  germanates  corresponding  to  silicates.  Tin  is  a  still  more 
basic  element,  forming  well-marked  salts  with  acids  ;  but  it  is  also 
acidic,,  and  with  alkalies  forms  stannates. 

Carbon  and  silicon  exhibit  a  close  relationship.  They  both 
form  allotropes,  which  correspond  in  many  respects.  They  both 
unite  with  hydrogen,  forming  the  analogous  compounds  CH4  and 
SiH4  ;  and  with  hydrogen  and  chlorine  they  form  the  similarly  con- 
stituted compounds,  chloroform,  CHC13  ;  and  silicon  chloroform, 
SiHCl3. 

Tin  and  lead  approach  more  nearly  to  each  other,  especially  in 
their  physical  properties,  than  to  the  other  members  of  the  family. 
They  both  form  compounds,  in  which  the  metals  function  both 
as  divalent  and  tetravalent  elements  ;  although  in  the  case  of 
lead  (as  often  happens  with  the  heaviest  metals  of  a  family),  the 
element  exhibits  much  greater  readiness  to  act  in  the  lower  state 
of  atomicity.  Until  quite  recently  (1893)  no  compound  was  known 
in  which  an  atom  of  lead  is  united  with  four  monovalent  atoms, 
although  lead  ethide,  Pb(C2H5)4,  had  been  obtained.  Now,  how- 
ever, the  compound  PbCl4  has  been  produced,  corresponding  to 
SnCl4,  which  it  resembles  in  many  respects  ;  and  still  more 
recently  (1894)  the  tetrafluoride  has  been  obtained. 

Carbon,  as  usual  with  the  typical  elements,  stands  apart  from 
the  other  members  of  the  family  in  many  of  its  attributes.  Thus, 
its  oxides  are  both  gaseous  ;  it  also  forms  a  vast  number  of  com- 
pounds with  hydrogen,  oxygen,  and  nitrogen,  the  study  of  which 
constitutes  the  science  of  organic  chemistry.  This  element  has 
already  been  treated  in  Part  II.  (page  285). 


SILICON. 

Symbol,  Si.     Atomic  weight =28. 3. 

Occurrence. — Silicon  is  not  known  to  occur  in  the  uncombined 
state,  although  in  combination  it  is  the  most  abundant  and  widely 
distributed  of  all  the  elements,  with  the  exception  of  oxygen.  In 
combination  with  oxygen,  as  silicon  dioxide  or  silica,  SiO2,  it 
occurs  &s>  flint,  sand,  quartz,  rock  crystal,  and  chalcedony /  while 
in  combination  with  oxygen  and  such  metals  as  calcium,  magnesium, 
and  aluminium,  it  occurs  in  clay  and  soil,  and  constitutes  a  large 
number  of  the  rocks  which  make  up  the  earth's  crust.  Silicon,  in 


Silicon  629 

combination  with  oxygen,  is  also  met  with  in  the  vegetable  kingdom, 
being  absorbed  by  plants  from  the  soil. 

Modes  of  Formation.  —  (i.)  Silicon  maybe  obtained  by  strongly 
heating  a  mixture  of  potassium  silico-fluoride  and  potassium  — 


The  mass,  after  cooling,  is  treated  with  water,  which  dissolves 
the  potassium  fluoride,  leaving  the  liberated  silicon. 

(2.)  This  element  may  also  be  prepared  by  heating  sodium  in  a 
stream  of  the  vapour  of  silicon  tetrachloride  — 


(3.)  In  an  impure  state,  mixed  with  magnesium  silicide,  it  may 
also  be  obtained  by  heating  a  mixture  of  dry  white  sand  with 
about  four  times  its  weight  of  dry  magnesium  powder  in  a  hard 
glass  tube. 

As  obtained  by  either  of  these  methods  the  silicon  is  in  the  form 
of  an  amorphous,  dark-brown  powder. 

(4.)  Silicon  is  obtained  in  a  crystalline  condition  by  passing  a 
slow  stream  of  the  vapour  of  silicon  tetrachloride  over  aluminium, 
previously  melted  in  a  current  of  hydrogen  ;  the  volatile  aluminium 
chloride  passes  on  in  the  stream  of  gas,  and  the  liberated  silicon 
dissolves  in  the  excess  of  aluminium  — 

3SiCl4  +  4A1  =  3Si  +  2A12C16. 

As  the  mass  cools,  silicon  is  deposited  in  the  form  of  long,  lustrous, 
needle-shaped  crystals. 

(5.)  The  most  convenient  method  for  the  preparation  of  crystal- 
lised silicon  consists  in  heating  in  a  crucible  a  mixture  of  3  parts 
of  potassium  silico-fluoride,  I  part  of  sodium,  and  4  parts  of  granu- 
lated zinc.  The  regulus  so  obtained  contains  crystallised  silicon. 
It  is  gently  heated,  and  the  excess  of  zinc  drained  away,  the  re- 
mainder being  removed  by  treatment  with  acids. 

Properties.—  Amorphous  Silicon,  as  obtained  by  the  reactions 
Nos.  i  and  2,  is  a  dark-brown  amorphous  powder,  having  a  specific 
gravity  of  2.15.  When  heated  in  the  air  it  burns  with  the  forma- 
tion of  silicon  dioxide,  which,  being  non-  volatile,  coats  the  particles 
of  the  element  and  protects  it  from  complete  oxidation.  It  burns 
when  heated  in  a  stream  of  chlorine,  with  formation  of  silicon 
tetrachloride.  It  is  insoluble  in  water,  and  in  all  acids  except 


630  Inorganic  Chemistry 

hydrofluoric  acid,   in  which  it    dissolves,  with  the   formation    of 
silico-fluoric  (or  hydrofluosilicic)  acid  and  evolution  of  hydrogen  — 

Si  +  6HF  =  H2SiF6  +  2H2. 

On  boiling  with  potassium  hydroxide  it  forms  potassium  silicate 
and  hydrogen  — 


Crystallised  Silicon.—  As  obtained  by  reactions  Nos.  4  and  5, 
silicon  is  a  brilliant,  steely-grey  substance,  crystallised  in  needles 
derived  from  the  orthorhombic  pyramid.  The  specific  gravity  of 
the  crystals  is  2.34  to  2.49.  Crystallised  silicon  does  not  burn  in 
oxygen,  even  when  strongly  heated  ;  it  burns  when  heated  in 
chlorine,  and  takes  fire  spontaneously  when  brought  into  fluorine. 
It  is  not  soluble  in  any  acid  except  a  mixture  of  nitric  and  hydro- 
fluoric acids.  Crystallised  silicon  is  very  hard,  being  capable  of 
scratching  glass.  When  silicon  is  exposed  to  a  high  temperature, 
out  of  contact  with  air,  it  becomes  denser  and  harder,  and  has 
been  obtained  in  the  form  of  small,  steel-grey  nodules,  showing  a 
crystalline  structure,  and  having  a  specific  gravity  as  high  as  3.0.* 

Silicon  Hydride,  SiH4.  —  This  compound  is  evolved  at  the 
negative  electrode  (along  with  hydrogen)  when  dilute  sulphuric 
acid  is  electrolysed,  the  electrodes  consisting  of  aluminium  con- 
taining silicon. 

In  an  impure  condition,  also  mixed  with  hydrogen,  this  gas  may 

*  Although  silicon  in  combination  is  such  an  abundant  element,  constituting, 
as  it  does,  about  one-fourth  of  the  total  weight  of  the  solid  crust  of  the  earth, 
in  the  free  state  it  must  still  be  regarded  as  somewhat  of  a  rarity,  and  con- 
sequently a  good  deal  of  uncertainty  exists  as  to  its  properties.  From  differ- 
ences that  have  been  observed  in  the  substance,  as  obtained  by  different 
methods,  and  from  the  close  analogy  that  exists  between  silicon  and  carbon, 
it  was  at  one  time  believed  that  three  allotropes  of  this  element  existed,  corre- 
sponding to  those  of  carbon.  Amorphous  silicon  was  considered  to  represent 
charcoal.  A  crystalline  substance  obtained  by  Wohler,  by  heating  potassium 
silico-fluoride  with  aluminium,  has  been  regarded  as  corresponding  to  graphite, 
and  called  graphitic  silicon  ;  while  the  octahedral  crystals  of  silicon  prepared 
by  reactions  4  and  5  given  above  (Deville)  were  thought  to  be  the  analogue  of 
diamond  ;  and  this  substance  has,  therefore,  been  called  diamond  or  adaman- 
toid  silicon.  There  is  considerable  doubt  as  to  whether  the  silicon  obtained 
by  all  these  various  methods  was  sufficiently  pure  to  warrant  this  classification, 
and  this  doubt  is  not  diminished  by  the  recently  discovered  fact  that  silicon 
unites  with  carbon,  forming  a  hard  crystalline  substance  which  has  received 
the  name  carborundum. 


Liquid  Silicon  Hydride  631 

be  obtained  by  the  action  of  hydrochloric  acid  upon  magnesium 
silicide  — 

SiMg2+4HCl  = 


(Magnesium  silicide  for  this  reaction  may  be  prepared  by  fusing 
together  in  a  covered  crucible  a  mixture  of  dry  magnesium  chloride 
40  parts,  dry  sodium  chloride  10  parts,  sodium  silico-fluoride  35 
parts,  and  metallic  sodium  20  parts.) 

Pure  silicon  hydride  is  prepared  by  acting  upon  triethyl  silico- 
formate  with  metallic  sodium.  The  mode  of  action  of  the  sodium 
is  not  known  ;  the  ethyl  silico-formate  breaks  up  into  silicon  hydride 
and  ethyl  silicate  — 


4SiH(OC2H6)3  =  SiH4  +  3Si(OC2H5)4. 

Properties.  —  Silicon  hydride  is  a  colourless  gas.  As  obtained 
by  the  first  two  methods  it  inflames  spontaneously.  The  pure 
gas  does  not  possess  this  property.  Its  ignition-point,  however, 
is  very  low,  and  if  the  gas  be  slightly  warmed,  or  if  a  jet  of  it  be 
caused  to  impinge  upon  an  object  a  few  degrees  above  the  ordinary 
temperature,  the  gas  at  once  takes  fire  and  burns  with  a  brightly 
luminous  flame  :  it  is  also  rendered  spontaneously  inflammable 
by  reduction  of  pressure  or  by  admixture  with  hydrogen.  When 
brought  into  chlorine  the  gas  takes  fire,  with  formation  of  silicon 
chloride  and  hydrochloric  acid. 

When  treated  with  an  aqueous  solution  of  sodium  or  potassium 
hydroxide,  silicon  hydride  is  decomposed,  giving  the  alkaline 
silicate  and  evolving  four  times  its  own  volume  of  hydrogen  — 

SiH4  +  2NaHO  +  H2O  =  SiO(NaO)2  +  4H2. 

Liquid  Silicon  Hydride,  Si2H6.—  This  compound,  which  has 
quite  recently  been  discovered,*  is  obtained  by  passing  the  pro- 
ducts from  the  action  of  dilute  hydrochloric  acid  upon  magnesium 
silicide  through  a  vessel  cooled  by  liquid  air  or  oxygen,  and 
separating  the  condensed  products  by  fractionation. 

Properties.  —  Liquid  silicon  hydride  is  a  colourless  mobile  liquid 
boiling  at  +  52°.  It  may  be  frozen  by  means  of  liquid  air  to  a 
white  crystalline  solid,  melting  at  -  1  38°.  The  liquid  is  spontaneously 
inflammable  in  air  at  the  ordinary  temperature,  burning  with  a  bright 
white  flame  and  depositing  amorphous  silicon  and  silicon  dioxide. 

*  Moissan  and  Smiles,  Comptes  Rendus,  March  1902. 


632  Inorganic  Chemistry 

If  a  small  quantity  of  the  liquid  be  vaporised  into  an  atmosphere 
of  hydrogen,  the  hydrogen  acquires  the  property  of  spontaneous 
inflammability  in  contact  with  the  air.  Liquid  silicon  hydride  is 
immediately  attacked  by  an  aqueous  solution  of  potash,  yielding 
potassium  silicate  and  hydrogen  — 

Si2H6  +  4KHO  +  2H2O  =  2SiO(KO)2  +  7H2. 

Silicon  Fluoride,  SiF4.  —  This  compound  is  formed  when  silicon 
is  brought  into  fluorine,  the  silicon  taking  fire  spontaneously  in 
the  gas. 

It  is  prepared  by  the  action  of  sulphuric  acid  upon  a  mixture  of 
powdered  fluorspar  and  sand  — 


Properties.  —  Silicon  fluoride  is  a  colourless,  fuming  gas.  It  is 
not  inflammable,  and  does  not  support  combustion.  It  is  decom- 
posed by  water  into  hydrofluosilicic  acid  and  silicic  acid,  hence  the 
gas  cannot  be  collected  over  water  — 

3SiF4  +  3H2O  =  2H2SiF6+H2SiO3. 

The  silicic  acid  is  precipitated  as  a  gelatinous  mass.  Each 
bubble  of  gas  as  it  comes  in  contact  with  the  water  is  at  once 
decomposed,  and  a  little  sack-like  envelope  of  silicic  acid  is 
formed  round  it.  On  filtering  the  liquid,  a  solution  of  hydrofluo- 
silicic acid  is  obtained.  When  silicon  fluoride  is  passed  over 
strongly  heated  silicon,  a  white  powder  is  obtained  having  the 
composition  Si2F6. 

Silicon  Chloride,  SiCl4,  is  formed  when  silicon  is  heated  in  a 
stream  of  chlorine.  Under  these  circumstances  the  silicon  burns 
in  the  gas. 

It  is  obtained  by  heating  an  intimate  mixture  of  silica  and 
carbon  in  a  stream  of  chlorine,  and  passing  the  products  through 
a  cooled  tube  — 


Properties.  —  Silicon  chloride  is  a  colourless  liquid  which  fumes 
strongly  in  moist  air  and  boils  at  58.3°.  It  is  decomposed  by 
water  into  silicic  and  hydrochloric  acids  — 

=  Si(HO)4 


Silicon  Dioxide  633 

and  the  silicic  acid  so  formed  passes  either  entirely  or  in  part 
into  the  dibasic  acid,  thus  — 

Si(HO)4  =  SiO(HO)2+H,O. 

Disilicon  Hexachloride  (also  known  as  silicon  trichloride),  Si2Cl6,  is  formed 
when  the  vapour  of  silicon  tetrachloride  is  passed  over  strongly  heated 
silicon  — 

l4  +  Si=2Si2Cl6. 


It  may  be  prepared  by  gently  heating  the  corresponding  iodine  compound 
with  mercuric  chloride— 

Si2I6  +  3HgCl2  =  Si2Cl6  +  3HgI2. 

Properties.  —  Disilicon  hexachloride  is  a  mobile,  colourless,  fuming  liquid, 
which  boils  at  147°  and  crystallises  at  -  i°.  When  the  liquid  is  boiled  and  the 
hot  vapour  allowed  to  escape  into  the  air,  it  spontaneously  ignites. 

Silicon  forms  two  compounds  with  bromine  and  with  iodine  corresponding 
to  the  chlorides,  namely  — 

SiBr4;  Si2Br6  ;  SiI4  ;  Si2I6. 

Silicon  Dioxide,  SiO2,  occurs  in  nature  in  a  more  or  less  pure 
form  in  a  large  number  of  minerals,  some  of  which  have  already 
been  alluded  to  as  natural  compounds  of  silicon.  Silicon  dioxide 
in  an  amorphous  form  is  met  with  in  the  different  varieties  of 
opal,  and  in  enormous  quantities  in  the  deposit  known  as  kiesel- 
guhr.  This  substance  consists  of  the  remains  of  extinct  dia- 
tomaceae,  and  is  met  with  in  various  parts  of  Germany.  In  a 
crystalline  condition  silica  occurs  as  quartz  or  rock  crystal,  and 
also  in  a  rarer  form  as  tridymite. 

Modes  of  Formation.—  (  i.)  Silicon  dioxide  is  formed  when 
amorphous  silicon  is  burnt  in  air  or  oxygen. 

(2.)  It  may  be  prepared  by  heating  silicic  acid,  which  readily  parts 
with  water  and  leaves  pure  silicon  dioxide  as  a  light  white  amor- 
phous powder  — 

Si(HO)4=SiO2  +  2H2O  ;  or 
SiO(HO)2  =  SiO2  +  H2O. 

(3.)  In  minute  crystals,  silicon  dioxide  is  obtained  by  strongly 
heating  a  solution  of  an  alkaline  silicate  in  a  sealed  glass  tube, 
whereby  a  portion  of  the  silica  of  the  glass  is  dissolved.  When 
this  solution  is  cooled,  silicon  dioxide  is  deposited.  If  the  crys- 
tallisation takes  place  above  a  temperature  of  180°,  crystals  of 
quartz  are  obtained  ;  if  below  this  point,  it  deposits  crystals  of 


634 


Inorganic  Chemistry 


tridymite,  while  at  ordinary  temperatures  the  silica  is  deposited 
in  the  amorphous  condition.  Much  larger  quartz  crystals  have 
been  obtained  by  the  prolonged  heating  to  250°  of  a  10  per  cent, 
aqueous  solution  of  silicic  acid  (obtained  by  dialysis)  in  stout 
sealed  glass  flasks. 

Properties. — In  the  crystalline  condition  as  quartz,  silicon 
dioxide  forms  prismatic  crystals  belonging  to  the  hexagonal  sys- 
tem, terminating  in  hexa- 
gonal pyramids.  Fig.  148 
represents  a  mass  of  quartz 
or  rock  crystal. 

The  purest  forms  of  rock 
crystal  are  perfectly  colour- 
less, having  a  specific  gravity 
of  2.69,  and  are  sufficiently 
hard  to  cut  glass.  When  cut 
and  polished,  it  exhibits  a 
brilliancy  not  far  inferior  to 
that  of  the  diamond,  and  is 
occasionally  substituted  for 
this  gem. 

Quartz  is  often  found 
coloured  by  the  presence  of 
small  quantities  of  impurities, 
as  in  the  varieties  known  as 
amethyst  quartz  and  smoky 
quartz^  and  in  great  quanti- 
ties as  milky  quartz. 

The  variety  of  silicon  di- 
oxide known  as  tridymite  is 

found  as  minute  crystals  in  cavities  in  certain  specimens  of  trachytic 
rocks.  The  crystalline  form  of  tridymite,  although  belonging  to 
the  hexagonal  system,  is  distinct  from  that  of  quartz,  and  the  crystals 
are  frequently  met  with  grown  together  in  the  manner  known  as 
twin-crystals. 

Amorphous  silicon  dioxide,  as  it  occurs  in  nature,  is  a  translu- 
cent substance  having  a  conchoidal  or  vitreous  fracture  ;  its  specific 
gravity  is  2.3.  As  artificially  prepared,  it  is  a  soft  white  powder 
whose  specific  gravity  is  2.2.  At  the  temperature  of  the  oxy- 
hydrogen  flame,  silicon  dioxide  melts  to  a  transparent  glass-like 
substance  which  is  capable  of  being  drawn  out  into  fine  threads 


FIG.  148. 


Silicic  Acids  635 

resembling  spun  glass.  These  fibres  possess  many  valuable  pro- 
perties, and  are  employed  by  physicists  in  delicate  instruments  of 
precision. 

Silicon  dioxide  is  insoluble  in  water  and  in  all  acids  with  the 
exception  of  hydrofluoric  acid.  It  dissolves  in  alkalies,  and  the 
amorphous  powder  can  be  dissolved  in  a  boiling  solution  of  sodium 
carbonate.  Many  natural  hot  springs  contain  silica  held  in  solu- 
tion as  an  alkaline  silicate,  and  on  exposure  to  atmospheric  carbon 
dioxide  the  silicate  is  decomposed  with  the  deposition  of  silica  and 
the  reformation  of  an  alkaline  carbonate.  The  enormous  quantities 
of  siliceous  sinter  deposited  by  geysers  at  Rotomahama,  New  Zea- 
land, were  formed  in  this  way.  When  fused  with  sodium  carbo- 
nate, silicon  dioxide  is  converted  into  soluble  sodium  silicate — 

SiO2  +  2Na2CO3  =  2CO2+Si(NaO)4. 

Silieie  Aeids. — Silicon  dioxide  is  capable  of  forming  weak 
polybasic  acids,  but  from  the  readiness  with  which  they  give  up 
water  it  is  probable  that  none  have  ever  been  obtained  in  a  state 
of  purity.  The  compound  represented  by  the  formula  Si(HO)4  is 
known  as  orthosilicic  acid,  and  is  tetrabasic.  By  the  loss  of  one 
molecule  of  water  it  forms  metasilicic  acid,  SiO(HO)2.  When 
hydrochloric  acid  is  added  to  a  solution  of  an  alkaline  silicate,  a 
gelatinous  precipitate  is  obtained,  which  consists  of  the  dibasic 
acid  SiO(HO)2,  or  H2SiO3— 

SiO(NaO)2+2HCl  =  SiO(HO)2  +  2NaCl. 

If,  on  the  other  hand,  the  solution  of  alkaline  silicate  be  added 
cautiously  to  an  excess  of  hydrochloric  acid,  the  silicic  acid  remains 
in  solution,  and  is  probably  present  as  orthosilicic  acid,  Si(HO)4,  or 
H4Si04— 

SiO(NaO)2  +  2HCl  +  H2O  =  Si(HO)4  +  2NaCl. 

The  sodium  chloride  in  the  solution  may  be  removed  by  a  pro- 
cess of  separation  known  as  dialysis.  This  process,  discovered  by 
Graham,  is  based  upon  a  property  belonging  to  certain  classes  of 
substances,  of  passing  when  in  solution  through  certain  mem- 
branes. The  mixture  is  placed  in  an  apparatus  resembling  a 
small  tambourine  (Fig.  149)  (made  by  stretching  either  parch- 
ment or  parchment  paper  over  a  wooden  hoop),  which  is  then 
floated  upon  water.  The  sodium  chloride  passes  through  the 


636  Inorganic  Chemistry 

membrane,  while  the  silicic  acid  remains  behind  in  the  dialyser 
as  a  dilute  aqueous  solution.  Substances  in  solution  which  are 
capable  of  readily  diffusing  through  such  a  membrane  were  termed 
by  Graham  crystalloids /  while  others,  such  as  the  silicic  acid, 
which  either  do  not  pass  through  or  only  do  so  with  difficulty,  are 
known  as  colloids. 

This  aqueous  solution  of  silicic  acid  may  be  concentrated  by 
boiling,  and  further  by  evaporation  in  vacuo  over  sulphuric  acid, 
until  it  contains  about  21  per  cent,  of  tetrabasic  silicic  acid,  or  14 
per  cent,  of  silicon  dioxide.  In  this  condition  it  is  a  tasteless 
liquid,  having  a  feeble  acid  reaction.  It  cannot  be  preserved,  as 


FIG.  149. 

on  standing  it  solidifies  to  a  transparent  gelatinous  mass,  which 
has  approximately  the  composition  H2SiO3. 

Silicates. — The  silicates  constitute  a  large  class  of  important  minerals, 
many  of  which  are  of  extremely  complex  composition.  Some  of  the  simplest 
of  these  silicates  are  derived  from  the  dibasic  and  tetrabasic  acids  already 
described,  while  others  may  be  regarded  as  the  salts  of  a  number  of  hypo- 
thetical polybasic  silicic  acids,  derived  from  metasilicic  acid  by  the  gradual 
elimination  of  water.  Thus,  by  the  withdrawal  of  one  molecule  of  water  from 
two  molecules  of  metasilicic  acid,  an  acid  known  as  disilicic  acid  is  obtained, 
having  the  composition  Si2O3(HO)2)  or  2SiO2,H2O,  or  H2Si2O5— 

2SiO(HO)2=  H2O  +  Si2O3(HO)2. 

By  the  abstraction  of  one  molecule  of  water  from  two  molecules  of  ortho- 
silicic  acid  another  disilicic  acid  is  similarly  derived — 

2Si(HO)4=H20  +  Si20(HO)6,  or  2SiO2,3H2O,  or  H6Si2O7. 


Tin  637 

By  the  partial  withdrawal  of  water  from  three  molecules  of  silicic  acid  a 
number  of  hypothetical  trisilicic  acids  may  be  derived,  such  as — 

3SiO2,2HoO  or  H4Si3Os ;  3SiO2,5H2O  or  H10Si3On  ; 
3Si02,7H20  or  H14Si3O13. 

Silicates  derived  from  an  acid  containing  one  atom  of  silicon  are  termed 
monosilicates ;  those  from  acids  with  two  or  three  atoms  of  silicon  respec- 
tively, disilicates  and  trisilicates. 

Thus,  the  mineral  peridote  is  a  monosilicate,  Mg2SiO4. 

Serpentine  is  a  disilicate,  Mg3Si2O7,  and 

Felspar,  or  orthoclase,  is  a  trisilicate,  Al2K2(Si3O8}2. 


TIN. 

Symbol,  Sn.     Atomic  weight  =119. 

Occurrence.  —  Tin  does  not  occur  in  nature  in  the  uncombined 
state.*  It  is  met  with  chiefly  as  the  oxide  SnO2  in  the  mineral 
tin-stone  or  cassiteriteft  which  is  found  in  immense  deposits, 
although  in  comparatively  few  localities.  It  is  usually  associated 
with  arsenical  ores,  copper  pyrites,  wolfram  (a  tungstate  of  iron 
and  manganese),  and  other  minerals.  Occasionally  it  is  met  with 
in  nodules  of  nearly  pure  oxide,  known  as  stream-tin. 

Mode  Of  Formation.—  Tin  is  obtained  exclusively  from  tin- 
stone; and  the  process  with  ordinary  ore  consists  of  three  opera- 
tions, namely  —  (i)  calcining,  (2)  washing,  (3)  reducing  or  smelting. 
If  the  ore  be  nearly  pure  tin-stone  it  may  be  at  once  smelted. 

The  finely  crushed  ore,  after  being  washed  from  earthy  matters, 
is  calcined  in  a  reverberatory  furnace.  The  sulphur  and  arsenic 
pass  away  as  sulphur  dioxide  and  arsenious  oxide,  and  are  led  into 
condensing  flues,  where  the  arsenic  deposits  and  is  collected.  The 
iron  and  copper  are  oxidised  to  oxide  and  sulphate.  This  calcina- 
tion is  sometimes  conducted  in  the  revolving  calciner,  shown  on 
page  486.  The  calcined  ore  is  next  washed,  whereby  copper 
sulphate  is  dissolved,  and  the  iron  oxide  and  other  light  matters 
are  removed.  The  purified  ore  is  then  mixed  with  powdered 
anthracite  and  smelted  in  a  reverberatory  furnace  — 


*  Metallic  tin  has  been  found  in  Bolivia,  but  its  origin,  whether  natural  or 
artificial,  is  doubtful. 

f  Cassiterides,  the  ancient  name  for  the  British  Isles,  is  derived  from  the 
fact  that  tin-stone  was  found  in  large  quantities  in  Devonshire  and  Cornwall. 


638  Inorganic  Chemistry 

The  metal  so  obtained  is  purified  by  first  heating  it  upon  the 
hearth  of  a  similar  furnace  until  the  more  readily  fusible  tin  melts 
and  flows  away  from  the  associated  alloys  ;  and  afterwards  by 
stirring  into  the  molten  tin  so  separated  billets  of  green  wood, 
which  results  in  the  separation  of  a  scum  or  dross  carrying  with  it 
the  impurities. 

Properties.— Tin  is  a  bright  white  metal,  which  retains  its 
lustre  unimpaired  in  the  air.  It  is  sufficiently  soft  to  be  cut  with  a 
knife,  but  is  harder  than  lead,  although  less  hard  than  zinc.  At 
ordinary  temperatures  it  is  readily  beaten  out  into  leaf  (known  as 
tinfoil),  and  may  be  drawn  into  wire  ;  but  at  temperatures  a  little 
below  its  melting-point  (228°)  it  becomes  brittle  and  may  be 
powdered.  Tin  may  be  obtained  in  the  form  of  crystals  by  melt- 
ing a  quantity  of  the  metal  in  a  crucible,  and  when  partially 
solidified  pouring  out  the  remaining  liquid  portion.  Its  crystalline 
character  is  also  seen  by  pouring  over  the  surface  of  a  block  of 
cast  tin  or  a  sheet  of  ordinary  tinned  iron  a  quantity  of  warm 
dilute  aqua  regia,  when  the  surface  of  the  metal  will  immediately 
exhibit  a  beautiful  crystalline  appearance. 

When  a  bar  of  tin  is  bent  it  emits  a  faint  crackling  sound,  and 
if  quickly  bent  backwards  and  forwards  two  or  three  times  the 
metal  becomes  perceptibly  hot  at  the  point  of  flexure.  These 
phenomena  are  due  to  the  friction  of  the  crystalline  particles. 
Ordinary  tin  has  a  specific  gravity  about  7.2  ;  but  if  the  metal  be 
exposed  to  the  prolonged  influence  of  very  low  temperatures,  it 
loses  its  crystalline  character  and  appears  of  a  grey  colour.  In 
this  condition  its  specific  gravity  is  5.8  ;  and  it  is  believed  to  be 
an  allotropic  modification  of  the  element.  When  strongly  heated 
tin  takes  fire  and  burns,  forming  stannic  oxide,  SnO2.  It  is 
oxidised  by  both  sulphuric  and  nitric  acids  ;  thus,  when  heated 
with  strong  sulphuric  acid,  stannous  sulphate  and  sulphur  dioxide 
are  produced — 

Sn  +  2H2SO4  =  SnSO4+SO2  +  2H2O. 

The  strongest  nitric  acid  (specific  gravity,  1.5)  is  without  action 
upon  tin.  Ordinary  concentrated  nitric  acid  (specific  gravity,  1.24) 
attacks  it  with  violence,  forming  metastannic  acid  (page  640),  while 
in  cold  dilute  acid  it  slowly  dissolves  with  the  production  of  stannous 
nitrate — 

=  4Sn(NO3)2 


Stannous  Oxide  639 

The  ammonia  unites  with  another  portion  of  nitric  acid,  forming 
ammonium  nitrate.  Strong  hydrochloric  acid  converts  it  into 
stannous  chloride,  with  evolution  of  hydrogen. 

Tin  is  extensively  employed  in  the  process  of  tinning,  which 
consists  in  coating  other  metals  with  a  thin  film  of  tin  by  dipping 
into  a  bath  of  the  molten  metal.  Ordinary  tin-plate  (or  in  common 
parlance,  "  tin,"  the  material  of  which  articles  generally  called 
"  tins  "  are  made)  is  thin  sheet-iron  which  has  been  thus  super- 
ficially coated  with  tin. 

Alloys  Of  Tin.  —  Tin  enters  into  the  composition  of  a  large 
number  of  useful  alloys.  With  lead,  tin  will  mix  in  all  proportions, 
and  many  alloys  are  in  use  consisting  of  these  two  metals.  They 
are  all  white,  and  melt  at  temperatures  lower  than  that  of  either 
constituent. 

Pewter  contains  3  parts  of  tin  to  i  part  of  lead.  Common 
solder  consists  of  I  part  tin  and  I  part  lead,  while  coarse  and  fine 
solder  contain  half  and  twice  this  proportion  of  tin  respectively. 
With  copper,  the  most  important  alloys  are  the  various  brasses 
and  bronzes.  Britannia  metal  contains  tin  84  parts,  antimony  10 
parts,  copper  4  parts,  and  bismuth  2  parts.  Tin  is  a  constituent 
also  of  the  so-called  fusible  alloys  (see  Bismuth,  page  500). 

Oxides  of  Tin.  —  Two  oxides  are  definitely  known,  namely, 
stannous  oxide,  SnO,  and  stannic  oxide,  SnO2.  The  monoxide  is 
a  base,  yielding  the  stannous  salts;  the  dioxide  is  both  a  basic  and 
an  acidic  oxide. 

Stannous  Oxide,  SnO,  is  obtained  by  heating  stannous  oxalate 
out  of  contact  with  air,  thus  — 

SnC2O4=  SnO  +  CO2  +  CO. 

When  sodium  carbonate  and  stannous  chloride  are  mixed,  carbon 
dioxide  is  evolved,  and  the  white  hydrated  oxide  is  precipitated, 
thus— 


When  this  hydrated  oxide  is  boiled  with  insufficient  caustic 
alkali  to  dissolve  it,  the  undissolved  portion  is  dehydrated  and 
converted  into  the  black  monoxide. 

When  heated  in  the  air,  stannous  oxide  becomes  incandescent, 
burning  to  the  dioxide.  It  is  soluble  in  acids,  forming  stannous 
salts.  The  solution  of  stannous  oxide  in  sodium  hydroxide  is 


640  Inorganic  Chemistry 

used  by  the  calico-printer,  and  is  known  commercially  as  sodium 
stannite. 

Stannic  Oxide,  SnO2  (tin  dioxide),  is  the  chief  ore  of  tin.  It  is 
formed  where  the  metal  is  burnt  in  the  air,  but  is  most  readily  pre- 
pared by  igniting  metastannic  acid. 

It  is  a  white  amorphous  powder,  which  changes  to  yellow  and 
brown  on  heating,  but  returns  to  its  original  condition  on  cooling. 
When  strongly  heated  in  a  stream  of  gaseous  hydrochloric  acid,  it 
may  be  obtained  in  small  crystals,  identical  with  the  natural  com- 
pound. Stannic  oxide  is  unacted  upon  by  acids  or  alkalies,  but 
in  contact  with  fused  potassium  hydroxide  it  is  converted  into 
potassium  stannate. 

Stannic  Acid,  H2SnO3,  or  SnO2,H2O,  is  obtained  in  a  hydrated 
condition,  as  a  white  gelatinous  precipitate,  when  calcium  car- 
bonate is  added  to  stannic  chloride  in  insufficient  quantity  for 
complete  precipitation.  When  the  precipitate  is  dried  in  vacuo, 
it  has  the  composition  H2SnO3.  The  equation  representing  its 
formation  may  be  expressed  thus  — 


Stannic  acid  forms  a  number  of  salts,  of  which  sodium  and 
potassium  stannates  are  the  most  important  —  the  former  being 
extensively  employed  as  a  mordant  in  dyeing,  under  the  name  of 
preparing  salt.  The  salts  have  the  composition  Na2SnO3,3H2O, 
and  K2SnO3,3H2O  respectively,  and  are  both  soluble  in  water. 

Metastannie  Acid,  H10Sn5O15,  is  obtained  as  a  white  amorphous 
powder  when  tin  is  acted  upon  by  strong  nitric  acid  ;  the  reaction 
may  be  represented  thus  —  • 

5Sn  +  20HNO3  =  H10Sn5O15  +  5H2O-f20NO2. 

The  composition  of  the  compound  depends  upon  the  particular 
temperature  at  which  it  is  dried.  This  acid  is  sometimes  regarded 
as  a  polymer  of  stannic  acid,  which  may  be  expressed  by  the 
formula  5(H2SnO3)  ;  metastannic  acid,  however,  appears  to  be 
dibasic,  forming  salts  in  which  two  only  of  the  hydrogen  atoms 
are  replaced  ;  its  composition  may  therefore  be  conveniently  ex- 
pressed thus  — 

H2SnO3,4SnO2,4H2O,  or  H2Sn5On,4H2P. 


Stannous  Chloride  641 

Potassium  and  sodium  metastannates  are  the  best  known  salts, 
their  formulae  being  — 

K2SnO3,4SnO2,4H2O,  and  Na2SnO3,4SnO2,4H2O. 

Stannous  Chloride,  SnCl2,  is  obtained  by  dissolving  tin  in 
hydrochloric  acid,  and  evaporating  the  solution,  when  monosym- 
metric  prisms  separate  out,  having  the  composition  SnCl2,2H2O. 
When  dried  in  vacuo  they  become  anhydrous.  The  anhydrous 
chloride  is  directly  obtained  when  tin  filings  and  mercuric  chloride 
are  heated  together  — 


The  reduced  mercury  volatilises  and  leaves  the  chloride,  which 
at  a  higher  temperature  may  be  distilled. 

Stannous  chloride  dissolves  in  a  small  quantity  of  water,  but 
with  an  excess  of  water,  or  on  exposure  to  the  air,  an  oxychloride 
(or  basic  chloride)  is  precipitated,  with  simultaneous  elimination  of 
hydrochloric  acid,  thus  — 

2SnCl2  +  2H20  =  SnCl2,SnO,H2O  +  2HCl. 

The  composition  of  this  oxychloride  may  also  be  expressed  by 
either  of  the  following  formulas  — 

Sn2OCl2,H2O,  or  2Sn(OH)Cl,  or  2(SnO,HCl). 

Stannous  chloride  is  a  powerful  reducing  agent,  as  it  readily 
combines  with  either  oxygen  or  chlorine  ;  thus,  when  added  to  a 
solution  of  mercuric  chloride,  the  latter  is  first  reduced  to  mer- 
curous  chloride,  which,  on  being  gently  warmed,  is  reduced  to 
metallic  mercury  — 

2HgCl2  +  SnCl2  =  Hg2Cl2+  SnCl4. 
Hg2Cl2  +  SnCl2=2Hg+  SnCl4. 

By  the  absorption  of  oxygen,  the  above  oxychloride  and  stannic 
chloride  are  formed,  thus  — 

3SnCl2+  O  +  H2O  =  SnCl2,SnO,H2O  +  SnCl4. 

Stannous  chloride  boils  at  a  temperature  about  606°.  The 
density  of  the  vapour  only  agrees  with  the  formula  SnCl2  at  tem- 
peratures above  900°,  at  lower  temperatures  its  vapour-density 
approaches  more  nearly  to  that  required  by  the  formula  Sn2Q4. 

2  s 


642  Inorganic  Chemistry 

Stannic  Chloride,  SnCl4)  is  obtained  by  passing  a  stream  of 
dry  chlorine  over  melted  tin  in  a  glass  retort,  or  by  heating  a 
mixture  of  powdered  tin  with  an  excess  of  mercuric  chloride,  when 
the  anhydrous  chloride  distils  over  as  a  colourless,  mobile,  fuming 
liquid,  which  boils  at  113.9°.  It  unites  with  water  with  evolution 
of  heat,  forming  hydrated  compounds  of  the  composition  SnCl4, 
3H2O  ;  SnCl4,5H2O,  and  SnCl4,8H2O.  The  compound  containing 
5H2O  is  employed  as  a  mordant,  and  is  commercially  known  as 
oxymuriate  of  tin. 

Stannic  chloride  combines  with  alkaline  chlorides,  forming 
double  chlorides  (sometimes  called  chloro-stannates\  such  as 
SnCl4,2NH4Cl,  and  SnCl4,2KCl. 

Stannous  Sulphide,  SnS.—  When  tinfoil  is  introduced  into 
sulphur  vapour  the  metal  takes  fire,  and  yields  a  leaden-coloured 
mass  of  stannous  sulphide. 

In  the  hydrated  condition  stannous  sulphide  is  precipitated  as 
a  brown  powder  when  sulphuretted  hydrogen  is  passed  through 
stannous  chloride  ;  on  drying,  this  becomes  black  and  anhydrous. 

Stannous  sulphide  dissolves  in  hot  concentrated  hydrochloric 
acid.  It  is  also  soluble  in  alkaline  polysulphides,  forming  thio- 
stannates,  thus  —  • 


SnS2  +  K2S  =  K2SnS3. 

On  the  addition  of  hydrochloric  acid  to  the  solution,  stannic 
sulphide  is  precipitated  — 


Stannic  Sulphide,  SnS2.  —  This  compound  cannot  be  formed 
by  heating  tin  and  sulphur  alone,  as  the  heat  of  the  reaction  is 
greater  than  that  at  which  stannic  sulphide  is  resolved  into 
stannous  sulphide  and  sulphur.  It  is  obtained  by  heating  tin 
amalgam,  sulphur  and  ammonium  chloride,  in  a  retort.  The  action 
that  takes  place  is  a  complicated  one,  various  products  being 
volatilised,  and  stannic  sulphide  remaining  in  the  retort  as  a  mass 
of  golden-yellow  scales.  Amongst  the  products  expelled  during 
the  process  are  ammonium  chloride,  sulphur,  mercuric  chloride, 
mercuric  sulphide,  and  sulphuretted  hydrogen.  The  ammonium 
chloride  present  probably  acts  by  the  formation  of  ammonium 
stannous  chloride,  as  an  intermediate  product,  which  is  then  de- 


Lead  643 

composed  with  the  production  of  stannic  sulphide  and  ammonium 
stannic  chloride  thus — 

2SnCl2,2NH4Ci  +  2S  =  SnS2  +  NH4Cl  +  SnCl4,2NH4Cl. 

Stannic  sulphide  is  a  golden  yellow  crystalline  substance  which, 
when  heated,  partially  sublimes  as  such,  but  is  for  the  most  part 
decomposed  into  the  monosulphide  and  free  sulphur.  It  is  largely 
used  as  a  pigment  known  as  mosaic  gold. 


LEAD. 

Symbol,  Pb.     Atomic  weight  =207.1. 

Occurrence.  —  Lead  has  been  found  in  small  quantities  in  the 
ancombined  state,  probably  reduced  from  its  ores  by  volcanic 
action. 

In  combination  with  sulphur  it  occurs  in  enormous  quantities  in 
the  mineral  galena,  PbS,  which  is  the  ore  from  which  the  metal  is 
chiefly  obtained.  Large  quantities  are  also  met  with  as  carbonate 
in  the  mineral  cerussite,  PbCO3.  Other  natural  compounds  are 
anglesite,  PbSO4  ;  lanarkite,  PbSO4,PbO  ;  matlockite,  PbCl2,PbO; 
pyromorphite,  3Pb3P2O8,PbCl2. 

Modes  of  Formation.  —  Lead  is  very  readily  reduced  from  its 
compounds,  and  on  this  account  was  one  of  the  earliest  known 
metals.  It  was  termed  by  the  Romans  plumbum  nigrum. 

Two  general  processes  are  made  use  of  for  the  reduction  of  lead 
from  its  ores  :  — 

In  the  first  method  (known  as  the  reduction  process}  the  lead 
sulphide  is  reduced  by  double  decomposition  with  lead  oxide  and 
sulphate,  which  are  formed  by  roasting  the  ore. 

In  the  second  (called  the  precipitation  process)  the  sulphide  is 
reduced  by  metallic  iron. 

(i.)  The  galena  is  introduced  into  a  reverberatory  furnace,  where 
it  is  partially  roasted,  whereby  a  portion  of  the  sulphide  is  oxidised 
to  sulphate  and  oxide  — 


2=PbSO4. 
O2  =  2PbO 

The  temperature  is  then  raised,  when  the  oxide  and  sulphate 


644 


Inorganic  Chemistry 


react  upon  a  further  portion  of  the  sulphide,  with  the  formation  of 
metallic  lead  and  the  evolution  of  sulphur  dioxide  — 


2PbO  +  PbS  =  3Pb+  SO,. 

This  method  of  lead  smelting  is  followed  when  the  ore  is  fairly 
free  from  other  metallic  sulphides.  The  reverberatory  furnace 
usually  employed  (known  as  the  Flintshire  furnace}  has  a  con- 
siderable depression,  or  well,  in  the  hearth,  where  the  metallic 


FIG.  150. 

lead  collects  during  the  process,  and  from  which  it  is  drawn  off 
into  a  metal  pot. 

The  same  process  is  carried  out  in  the  North  of  England,  and 
in  Scotland,  where  a  very  pure  lead  ore  is  employed,  upon  open 
shallow  hearths  (known  as  the  ore  hearth,  or  Scotch  hearth},  built 
under  a  brickwork  hood  or  chimney  in  such  a  manner  that  the 
fumes  of  lead  which  escape  are  caused  to  pass  into  condensing 
chambers.  Fig.  150  shows  such  a  hearth  in  section.  The  fire  of 
peat  and  coal  is  urged  by  a  small  blast  admitted  from  behind,  and 
the  ore  is  added  in  small  quantities  at  a  time.  The  reduced  metal, 
sinking  to  the  bottom,  runs  under  the  fire-bar  and  overflows 
from  the  shallow  hearth  down  a  channel  upon  an  inclined  stone 
surface  S  (called  the  ivork-stone\  into  an  iron  pot  P,  which  is  gently 


Lead  645 

heated  by  a  small  fire  to  enable  the  operator  to  ladle  the  metal  out 
into  moulds. 

(2.)  This  method  oflead  smelting  depends  upon  the  fact  that  at 
a  high  temperature  metallic  iron,  in  contact  with  lead  sulphide,  is 
converted  into  ferrous  sulphide,  with  separation  of  lead  — 


The  ores  (either  in  the  raw  state,  or  after  previous  calcination) 
are  smelted  in  a  blast-furnace  with  coke  and  either  metallic  iron 
or  such  materials  as  will  yield  iron  under  the  furnace  conditions. 
The  sulphide  of  iron,  along  with  other  metallic  sulphides,  rises  to 
the  top  of  the  molten  lead  as  a  matt  or  regulus,  while  above  this  a 
fusible  slag  collects,  consisting  chiefly  of  silicate  of  iron. 

The  lead  first  obtained  by  any  of  these  processes  usually  con- 
tains antimony,  tin,  copper,  and  other  metals.  These  impurities 
are  removed  by  heating  the  metal  in  a  shallow,  flat-bottomed 
reverberatory  furnace.  Most  of  the  admixed  metals  oxidise  before 
the  lead,  and  collect  in  the  dross  which  forms  upon  the  surface. 
This  process  is  known  as  the  softening  of  lead.  The  silver,  how- 
ever, which  is  always  present,  is  not  removed  by  this  operation, 
but  is  extracted  by  one  of  the  methods  for  desilverising  lead 
described  under  silver,  page  560. 

Properties.  —  Lead  is  a  soft,  bluish-white  metal,  which  when 
freshly  cut  exhibits  a  bright  metallic  lustre.  On  exposure  to  the 
air  its  bright  surface  becomes  quickly  covered  with  a  film  of  oxide. 
Lead  is  sufficiently  soft  to  be  scratched  with  the  finger  nail,  and 
it  leaves  a  black  streak  when  drawn  across  paper.  It  cannot  be 
hammered  into  foil  or  drawn  into  wire,  but  may  readily  be  obtained 
in  these  forms  by  rolling  and  pressing.  When  a  quantity  of 
melted  lead  is  allowed  partially  to  resolidify,  and  the  still  liquid 
portion  poured  off,  the  metal  is  obtained  in  the  form  of  octahedral 
crystals  belonging  to  the  regular  system.  Its  crystalline  nature  is 
also  readily  seen  by  submitting  a  solution  of  a  lead  salt  to  electro- 
lysis, when  the  metal  is  deposited  upon  the  negative  electrode  in 
beautiful  arborescent  crystals  with  a  brilliant  metallic  lustre  (Fig. 
151).  It  is  deposited  in  a  similar  form,  known  as  the  lead  tree,  by 
suspending  a  strip  of  zinc  in  such  a  solution.  The  specific  gravity 
oflead  is  11.3  ;  it  melts  at  330°  to  335°,  and  becomes  covered  with 
a  black  film  of  the  suboxide,  Pb2O  :  when  more  strongly  heated 
it  is  oxidised  to  the  monoxide,  PbO. 

Lead  is  rapidly  dissolved  by  nitric  acid,  but  hydrochloric  and 


646  Inorganic  Chemistry 

sulphuric  acids  are  almost  without  action  upon  it  in  the  cold.  Hot 
concentrated  hydrochloric  acid,  however,  slowly  converts  it  into 
lead  chloride. 

Lead  is  unacted  upon  by  pure  water  in  the  absence  of  air  ;  but 
m  contact  with  air  lead  hydroxide  is  formed,  which  is  slightly 
soluble  in  water.  By  the  action  of  atmospheric  carbon  dioxide  upon 
this  solution,  a  basic  carbonate  is  precipitated,  having  the  com- 
position 2PbCO3,Pb(HO)2.  The  solvent  action  of  water  upon  lead 
is  greatly  influenced  by  the  presence  of  various  dissolved  sub- 
stances in  the  water  ;  thus,  water  containing  small  quantities  of 


ammoniacal  salts,  notably  the  nitrate,  dissolves  lead  much  more 
rapidly,  and  the  same  is  the  case  with  water  charged  with  carbon 
dioxide  under  pressure.  In  the  latter  case  the  action  is  probably 
due  to  the  formation  of  a  soluble  acid  carbonate. 

Water,  on  the  other  hand,  containing  small  quantities  of  phos- 
phates and  carbonates,  especially  the  acid  calcium  carbonate,  are 
almost  entirely  without  action  upon  lead.  Certain  drinking  waters 
(such  as  the  Loch  Katrine  water),  which  on  account  of  their  purity 
exert  a  solvent  action  upon  the  lead  pipes  through  which  they  are 


Plumbic  Oxide  647 

conveyed,  are  rendered  incapable  of  acting  upon  the  lead  by  being 
first  filtered  through  chalk  or  animal  charcoal,  which  enables  them 
to  take  up  sufficient  calcium  carbonate  or  phosphate  to  prevent 
this  action. 

On  account  of  the  exhaustive  methods  of  desilverisation  to  which 
the  lead  is  subjected,  commercial  lead  possesses  a  degree  of  purity 
not  found  in  any  other  metal  as  commonly  met  with  ;  the  total 
amount  of  foreign  metals  present  in  ordinary  commercial  lead 
ranges  from  o.  I  to  0.006  per  cent. 

Lead  is  put  to  a  large  number  of  uses  in  the  arts,  on  account  of 
the  ease  with  which  it  can  be  worked,  and  its  power  of  resisting 
the  action  of  water  and  many  acids.  In  the  manufacture  of  lead 
pipes  advantage  is  taken  of  the  extreme  softness  of  the  metal  and 
the  readiness  with  which  it  can  be  pressed  into  shape  ;  the  lead, 
in  a  pasty  or  semi-molten  condition,  being  merely  squeezed,  or 
squirted,  through  a  steel  die  by  hydraulic  pressure. 

Lead  bullets  are  also  made  by  squeezing  the  metal  into  moulds  ; 
for  as  lead  contracts  on  solidification,  bullets  made  by  casting 
always  contain  a  small  cavity,  which  (unless  it  happens  to  form 
exactly  at  the  point  of  centre  of  gravity)  renders  the  flight  of  the 
bullet  untrue. 

Oxides  of  Lead.—  Five  oxides  of  lead  are  known,  having  the 
composition  Pb2O,  PbO,  Pb2O3,  Pb3O4,  PbO2. 

Lead  Suboxide  (plumbous  oxide],  Pb2O,  is  the  black  compound 
which  is  formed  when  lead  is  heated  to  its  melting-point.  It  is 
obtained  by  heating  plumbic  oxalate  to  about  300°  in  a  glass  tube 
or  retort  — 

2PbC204  =  CO  +  3CO2+  Pb2O. 

When  heated  in  the  air  it  burns,  forming  plumbic  oxide  ;  in  the 
absence  of  air  it  is  decomposed  into  the  same  oxide  and  metallic 
lead,  the  reactions  being  — 

Pb20  +  0  =  2PbO. 


In  contact  with  acids  it  decomposes  in  the  same  manner,  lead 
being  deposited,  and  the  plumbic  oxide  dissolving  in  the  acid  to 
form  a  plumbic  salt. 

Plumbic  Oxide  (lead  monoxide,  litharge,  massicot},  PbO,  is 
formed  when  lead  is  strongly  heated  in  the  air,  and  is  obtained  in 
large  quantities  in  the  cupellation  of  argentiferous  lead.  It  may 


648  Inorganic  Chemistry 

be  obtained  by  heating  lead  nitrate  or  carbonate,  and  it  is  produced 
when  any  of  the  other  oxides  are  heated. 

Plumbic  oxide  is  a  yellowish  powder,  known  commercially  as 
massicot^  which,  when  melted  and  resolidified,  is  obtained  as  a 
crystalline  mass,  known  as  litharge.  Plumbic  oxide  is  very  slightly 
soluble  in  water,  I  part  dissolving  in  7000  parts  of  water  :  this 
solution  is  alkaline,  and  on  exposure  to  the  air  absorbs  carbon 
dioxide,  forming  an  insoluble  basic  carbonate.  Plumbic  oxide  is 
dissolved  by  acids,  with  formation  of  the  salts  of  lead  ;  it  also 
dissolves  in  warm  potassium  or  sodium  hydroxide. 

This  oxide  forms  two  hydrated  compounds,  having  the  com- 
position 2PbO,H2O  and  3PbO,H2O.  The  former  is  obtained  as  a 
white  precipitate  when  ammonia  is  added  to  a  solution  of  lead 
acetate  ;  the  second,  by  the  action  of  ammonia  on  basic  lead 
acetate  at  25°. 

Lead  Sesquioxide,  Pb2O3,  is  obtained  as  an  orange-coloured 
precipitate  by  adding  sodium  hypochlorite  to  a  solution  of  plumbic 
oxide  in  potassium  hydroxide.  Heat  decomposes  it  into  oxygen 
and  plumbic  oxide.  Acids  convert  it  into  the  monoxide  and 
dioxide,  the  former  dissolving  and  yielding  a  salt  of  lead.  This 
oxide  may  be  regarded  as  a  compound  of  two  oxides,  PbO,PbO2. 

Triplumbie  Tetroxide  (red  lead,  minium],  Pb3O4,  is  obtained 
when  lead  carbonate,  or  monoxide,  is  subjected  to  prolonged 
heating  in  contact  with  air,  at  a  temperature  not  above  450°.  At 
higher  temperatures  it  again  gives  up  oxygen.  It  is  a  scarlet 
crystalline  powder,  varying  somewhat  in  colour,  according  to  its 
mode  of  preparation.  Dilute  acids  convert  it  into  PbO2  and 
2PbO,  the  latter  oxide  dissolving  to  yield  lead  salts.  With  strong 
hydrochloric  acid  and  sulphuric  acid  the  molecule  of  lead 
dioxide  is  acted  upon,  with  evolution  of  chlorine  and  oxygen 
respectively  — 

=  4H2O  +  3PbCl2+Cl2. 


When  red  lead  is  added  in  small  quantities  at  a  time  to  hot 
glacial  acetic  acid,  it  dissolves  entirely  in  the  acid,  and  the  liquid 
contains  lead  tetracetate  —  a  salt  which  is  of  interest  as  being  one  of 
the  few  salts  known  containing  tetravalent  lead  (see  page  628).  On 
cooling,  the  tetracetate  separates  out  as  pale  greenish-white  needles. 
The  salt  is  immediately  decomposed  by  water  ;  if  therefore  the 


Plumbic  Chloride  649 

acid  solution  of  it  be  poured  into  water,  a  brown  precipitate  of 
lead  peroxide  is  thrown  down,  and  acetic  acid  regenerated  — 


Red  lead*  is  employed  as  a  pigment,  and  also  in  the  manu- 
facture of  flint  glass. 

Plumbic  Peroxide  (lead  dioxide),  PbO2,  may  be  obtained  by 
the  action  of  dilute  nitric  acid  upon  red  lead  — 

Pb3O4  (or  PbO2,2PbO)  +  4HNO3  =  PbO2  +  2Pb(NO3)2  +  2H2O. 

Or  it  may  be  prepared  by  the  action  of  oxidising  agents  upon 
the  monoxide.  Thus,  when  chlorine  is  passed  through  an  alkaline 
solution,  in  which  the  monoxide  is  suspended,  or  when  bleaching- 
powder  is  added  to  a  solution  of  lead  acetate,  the  dioxide  is 
produced. 

The  dark-brown  deposit  which  forms  upon  the  positive  electrode 
when  a  solution  of  a  lead  salt  is  electrolysed,  also  consists  of  the 
dioxide. 

Plumbic  peroxide  is  a  brown  or  puce-coloured  powder.  It  is  a 
powerful  oxidising  substance,  and  when  gently  rubbed  with  flowers 
of  sulphur  in  a  warm  mortar  the  mass  suddenly  inflames.  When 
a  stream  of  sulphur  dioxide  is  passed  over  the  peroxide  in  a  tube, 
the  two  compounds  unite  to  form  lead  sulphate,  the  mass  becom- 
ing incandescent.  Nitric  acid  is  without  action  upon  it,  but 
hydrochloric  and  sulphuric  acids  act  upon  it  in  the  same  manner 
as  upon  red  lead.  When  strongly  heated  the  peroxide  gives  up 
oxygen,  and  is  converted  into  the  monoxide. 

When  plumbic  peroxide  is  boiled  with  strong  aqueous  potassium 
hydroxide  it  dissolves,  and  the  solution  deposits  crystals  of  potas- 
sium plumbate,  K2PbO3,3H2O.  This  compound  corresponds  with 
potassium  stannate,  K2SnO3,3H2O,  and  its  existence  shows  that 
lead  possesses,  although  to  a  very  feeble  extent,  the  acidic  properties 
exhibited  by  the  other  members  of  the  same  family  of  elements. 

Plumbic  Chloride  (lead  dichloride\  PbCLj,  is  obtained  as  a 
white  curdy  precipitate  when  hydrochloric  acid,  or  a  soluble 
chloride,  is  added  to  a  solution  of  a  lead  salt.  It  is  also  produced 

*  Commercial  red  lead  varies  considerably  in  composition,  and  although  it 
has  been  shown  that  a  definite  compound  exists  of  the  composition  Pb3O4 
(which  may  also  be  expressed  by  the  formula  2PbO,PbO2),  it  is  still  uncer- 
tain whether  there  are  not  other  compounds  consisting  of  these  two  oxides 
united  in  different  proportions. 


650  Inorganic  Chemistry 

by  the  action  of  boiling  hydrochloric  acid  upon  lead  in  the  pre- 
sence of  air-  It  is  best  prepared  by  dissolving  lead  oxide  or 
carbonate  in  hot  hydrochloric  acid,  when  the  lead  chloride  sepa- 
rates out  on  cooling  in  long  white,  lustrous,  needle-shaped  crystals 
belonging  to  the  rhombic  system.  Lead  chloride  is  soluble  in 
boiling  water  to  the  extent  of  about  4  parts  in  100  parts  of  water. 
On  cooling  the  solution  the  greater  part  of  the  salt  separates  out, 
and  at  o°  the  liquid  contains  0.8  part  in  solution.  The  presence 
of  hydrochloric  acid  and  soluble  chlorides  diminishes  the  solu- 
bility of  lead  chloride. 

When  heated  in  contact  with  air  it  is  converted  into  an  oxy- 
chloride,  of  the  composition  Pb2OCl2,  or  PbCl2,PbO,  corresponding 
with  the  natural  compound  matlockite.  This  compound,  in  the 
hydrated  condition,  Pb2OCl2,H2O,  is  prepared  on  a  large  scale  by 
the  addition  of  lime-water  to  a  solution  of  lead  chloride,  and  is 
employed  as  a  white  pigment,  known  as  Pattinsorts  white  lead. 

Cassel  yellow  is  an  oxychloride  of  lead  of  the  composition 
PbCl2,7PbO,  obtained  by  heating  lead  oxide  and  ammonium 
chloride. 

Lead  Tetraehloride  (fead perckloride\  PbCl4. — When  plumbic 
peroxide  is  dissolved  in  cold  concentrated  hydrochloric  acid,  a 
yellow  liquid  is  obtained,  which,  on  warming,  yields  chlorine, 
with  precipitation  of  lead  dichloride.  This  liquid  contains  the 
tetrachloride  of  lead  in  solution. 

When  lead  dichloride  is  suspended  in  hydrochloric  acid,  and 
chlorine  is  passed  through  the  mixture,  a  solution  of  lead  tetra- 
chloride is  obtained  ;  and  on  the  addition  of  ammonium  chloride, 
ammonium  plumbic  chloride,  PbCl4,2NH4Cl  (corresponding  to 
ammonium  stannic  chloride),  separates  out.  When  this  compound 
is  acted  upon  with  strong  sulphuric  acid,  in  the  cold,  lead  tetra- 
chloride separates  out  as  a  yellow  oily  liquid. 

Lead  tetrachloride  is  a  yellow,  highly-refracting,  fuming  liquid, 
which  decomposes  in  contact  with  moisture  into  lead  dichloride 
and  chlorine.  It  may  be  preserved  beneath  concentrated  sul- 
phuric acid.  With  small  quantities  of  water  it  forms  a  hydrated 
compound,  but  excess  of  water  decomposes  it  into  hydrochloric 
acid  and  lead  peroxide— 

PbCl4  +  2H2O  =  PbO2  +  4H  Cl. 

When  heated  with  strong  sulphuric  acid  to  about  105°,  it  suddenly 
decomposes  with  explosion. 


Lead  Carbonate  651 

Lead  Nitrate,  Pb(NO3)2,  is  obtained  by  dissolving  litharge  in 
nitric  acid.  The  salt  is  deposited  from  the  solution  in  the  form  of 
regular  octahedral  crystals.  It  is  soluble  in  water  to  the  extent  of 
50  parts  in  100  parts  of  water  at  the  ordinary  temperature.  When 
heated  it  evolves  nitrogen  peroxide  and  oxygen,  leaving  plumbic 
oxide  (page  242). 

On  boiling  an  aqueous  solution  of  lead  nitrate  with  lead  oxide, 
the  latter  dissolves,  and  the  solution  on  cooling  deposits  crystals 
of  a  basic  nitrate,  Pb(NO3)HO  or  Pb(NO3)2,PbO,H2O.  By  the 
addition  of  ammonia  to  a  solution  of  lead  nitrate,  other  basic 
nitrates  are  obtained,  which  may  be  regarded  as  consisting  of 
compounds  of  Pb(NO3)HO  with  PbO,  or  of  Pb(NO3)2  with  PbO 
and  H2O  in  varying  proportions. 

Lead  Carbonate,  PbCO3,  is  obtained  as  a  white  crystalline 
powder  by  the  addition 
of  ammonium  sesquicar- 
bonate  to  a  solution  of 
lead  nitrate.  It  occurs  in 
the  form  of  transparent 
rhombic  crystals  in  the 
mineral  ccrussite,  isomor- 
phous  with  arragonite. 
Lead  carbonate  is  almost 
insoluble  in  water,  but  is 
appreciably  dissolved  in 
water  charged  with  carbon  pIG  re2> 

dioxide.     When  sodium  or 

potassium  carbonate  is  added  to  a  solution  of  lead  nitrate,  basic 
carbonates  of  lead  are  precipitated,  varying  in  composition  with  the 
conditions  of  temperature.  The  most  important  of  the  basic  car- 
bonates is  white  lead,  2PbCO^Pb(YLO).^  This  compound  is  manu- 
factured on  a  large  scale  by  several  processes  for  use  as  a  pigment. 
The  oldest  process,  and  that  which  yields  the  best  product,  is  known 
as  the  Dutch  method.  It  depends  upon  the  action  of  acetic  acid 
upon  metallic  lead,  in  the  presence  of  moist  air  and  carbon 
dioxide.  The  lead,  cast  into  rough  gratings  in  order  to  expose 
a  large  surface,  is  placed  in  earthenware  pots,  as  shown  in  Fig. 
152.  A  small  quantity  of  dilute  acetic  acid  (in  the  old  Dutch 
process,  vinegar}  is  placed  in  the  pots,  and  the  gratings  of  lead, 
which  rest  upon  the  shoulder  of  the  pot,  are  piled  one  upon  the 
other.  These  pots  are  then  placed  upon  a  thick  bed  of  spent  tan- 


652  Inorganic  Chemistry 

bark  (in  the  original  method,  dung},  upon  the  floor  of  a  shed, 
and  covered  with  planks.  Upon  these  another  layer  of  tan- 
bark  is  spread,  and  a  second  row  of  pots  similarly  charged.  In 
this  manner  the  layers  of  pots  are  built  up  to  the  roof  of 
the  shed,  and  the  whole  allowed  to  remain  for  about  three 
months.  Such  a  stack  will  contain  many  tons  of  lead,  and 
about  65  gallons  of  dilute  acetic  acid  to  the  ton  of  metal. 
The  acid  is  gradually  vaporised  by  the  heat  developed  by  the 
fermenting  tan-bark,  which  results  first  in  the  formation  of  a 
basic  lead  acetate — 

2H(C2H302)  +  2Pb  +  O2=Pb(C2H3O2)2,Pb(HO)2. 

This  basic  acetate  is  then  acted  upon  by  the  carbon  dioxide 
evolved  during  the  fermentation,  with  the  production  of  a  mixture 
of  normal  lead  acetate  and  basic  lead  carbonate,  thus — 

3{Pb(C2H302)2)Pb(HO)2}  +  2C02=3Pb(CoH302)2  +  2PbC03)Pb(HO)2  +  2H20. 

And  the  lead  acetate,  in  the  presence  of  air  and  moisture,  reacts 
upon  a  further  portion  of  the  metal,  regenerating  the  basic  acetate, 
which  is  once  more  decomposed  by  carbon  dioxide — 

Pb(C2H302),  +  Pb  +  0  +  H2O  =  {Pb(C2H302)2,Pb(HO)2], 

In  this  cycle  of  reactions,  therefore,  the  acetic  acid  acts  as  a 
carrier,  a  comparatively  small  quantity  being  able  to  convert  an 
indefinite  amount  of  lead  into  'white  lead. 

White  lead  is  also  prepared  by  passing  carbon  dioxide  into  a 
solution  of  the  basic  acetate,  obtained  by  boiling  plumbic  oxide 
(litharge)  with  lead  acetate.  The  product,  however,  is  not  so 
opaque  as  that  obtained  by  the  former  method,  and  is  therefore 
not  so  valuable  as  a  pigment.  (This  method  is  known  as  the 
ClicJiy^  or  Thenard's  process.) 

Milner*s  process  consists  in  grinding  together  litharge,  sodium 
chloride,  and  water,  whereby  a  mixture  of  an  oxychloride  of  lead 
and  sodium  hydroxide  is  formed — 

5H2O  =  PbCl2,3PbO,4H2O  +  2 


Lead  Sulphate  653 

and  then  passing  carbon  dioxide  into  the  mixture,  which  converts 
it  into  'white  lead  and  sodium  chloride,  thus  —  • 


3[PbCl2,3PbO,4H2 

+  4[2PbCO3,Pb(HO)2]  +  11H2O. 

White  lead  is  a  heavy,  amorphous  powder,  whose  value  as  a 
pigment,  or  body  colour,  depends  upon  its  opacity  and  density. 
Although  this  compound  labours  under  the  disadvantages  of  being 
extremely  poisonous,  and  of  becoming  blackened  by  sulphuretted 
hydrogen,  no  substitute  for  it  has  yet  been  found  which  possesses 
the  same  "body  "  or  covering  power. 

Lead  Sulphate,  PbSO4.—  The  mineral  anglesite,  PbSO4,  occurs 
in  the  form  of  rhombic  crystals,  isomorphous  with  strontium  and 
barium  sulphates.  Lead  sulphate  is  obtained  as  a  white  powder} 
by  precipitating  a  lead  salt  with  sulphuric  acid  or  a  soluble 
sulphate.  It  is  soluble  in  water  only  to  an  extremely  slight  extent, 
and  still  less  in  dilute  sulphuric  acid,  but  strong  sulphuric  acid 
dissolves  it  readily.  It  also  dissolves  in  potassium  hydroxide,  and 
in  many  ammoniacal  salts,  notably  the  acetate,  and  in  sodium 
thiosulphate. 

An  acid  sulphate,  of  the  composition  PbSO4,H2SO4,H2O,  is 
obtained  by  boiling  the  normal  sulphate  with  sulphuric  acid  ;  and 
a  basic  sulphate,  PbSO4,PbO,  is  formed  by  the  action  of  ammonia 
upon  the  normal  salt. 

Lead  Disulphate,  Pb(SO4)2.—  This  substance  is  obtained  by  the 
electrolysis  of  sulphuric  acid  of  sp.  gr.  1.7  to  1.8  at  a  temperature 
not  above  30°,  employing  an  anode  of  lead.  The  cell  is  divided 
by  a  porous  partition,  and  the  compound  collects  as  a  muddy 
deposit  in  the  anode  compartment.  Lead  disulphate  is  a  cry- 
stalline substance  having  a  faint  greenish  colour.  It  is  immediately 
decomposed  by  water  into  lead  peroxide  and  sulphuric  acid  — 

Pb(SO4)2 


Sulphuric  acid  of  a  sp.  gr.  less  than  1.65  decomposes  it  in  a 
similar  manner,  but  in  concentrated  sulphuric  acid  it  is  slightly 
soluble,  loo  c.c.  acid  at  30°  dissolving  0.345  gram  of  the  com- 


654  Inorganic  Chemistry 

pound.  Concentrated  hydrochloric  acid  and  glacial  acetic  acid 
convert  it  respectively  into  lead  tetrachloride,  PbCl4,  and  lead 
tetracetate,  Pb(C2H3O2)4.  Each  of  these  compounds,  like  the 
disulphate,  is  decomposed  by  water  into  lead  peroxide  and 
the  respective  acid.  Double  salts,  such  as  K2Pb(SO4)3  and 
(NH4)2Pb(SO4)3,  have  been  prepared,  which  are  more  stable 
than  the  disulphate  itself. 

Lead  Sulphide,  PbS. — The  natural  sulphide,  galena,  is  found 
in  the  form  of  cubical  crystals,  possessing  very  much  the  colour 
and  the  metallic  lustre  of  freshly  cut  lead.  It  is  artificially  formed 
when  lead  is  heated  in  sulphur  vapour,  or  when  sulphuretted 
hydrogen  is  passed  through  a  solution  of  a  lead  salt. 

When  heated  in  vacuo,  or  in  a  stream  of  an  inert  gas,  lead 
sulphide  melts,  and  sublimes  in  the  form  of  small  cubes.  When 
heated  with  free  access  of  air  it  is  converted  into  lead  sulphate. 

Boiling  dilute  nitric  acid  converts  lead  sulphide  into  the  nitrate, 
with  separation  of  sulphur;  but  strong  nitric  acid  oxidises  it  into 
lead  sulphate.  It  is  decomposed  by  hot  concentrated  hydrochloric 
acid,  with  evolution  of  sulphuretted  hydrogen. 

When  sulphuretted  hydrogen  is  passed  into  a  solution  of  lead 
chloride,  the  precipitate  which  forms  is  first  yellow,  then  reddish- 
brown,  and  finally  black  ;  the  yellow  and  red  precipitates  are  com- 
pounds of  lead  chloride  and  lead  sulphide,  termed  sulphochlorides, 
having  the  composition,  PbS,PbCl2,  and  3PbS,PbCl2. 

The  compounds  of  lead  are  powerful  poisons,  and  when  con- 
tinuously taken  into  the  system  in  small  quantities,  they  act  as 
cumulative  poisons.  Painters  and  others  who  constantly  handle 
white  lead  are  liable  to  suffer  from  chronic  lead  poisoning. 


CHAPTER.   X 

ELEMENTS   OF   GROUP    V.  (FAMILY  A.} 

Vanadium,  V  =  5i.2;  Columbium,  05=93.5;    Tantalum,  Ta  =  i8i. 

THE  three  rare  metals  comprising  this  family  are  closely  related  to  each 
other,  and  also  to  the  elements  of  family  B  of  the  same  group,  namely,  the 
nitrogen  and  phosphorus  series. 

Vanadium  occurs  in  a  few  rare  minerals,  as  vanadite,  3Pb3(VO4)2,PbCl2 
(the  vanadium  analogue  of  pyromorphite)  ;  pucherite,  BiVO4  ;  mottramite, 
(PbCu)3(VO4)2,2(PbCu)(HO)2.  Small  quantities  also  occur  in  certain  iron 
ores,  the  vanadium  ultimately  finding  its  way  into  the  Bessemer  slag,  in 
which  it  has  been  found  concentrated  to  the  extent  of  1.5  per  cent. 

Metallic  vanadium  was  first  isolated  by  Roscoe  (1867),  although  its  existence 
was  previously  discovered  by  Del  Rio  (1801).  The  metal  is  extremely  difficult 
to  obtain,  as  at  a  red  heat  it  combines  with  oxygen  with  great  readiness, 
yielding  the  pentoxide  V2O5,  and  also  with  nitrogen,  forming  the  nitride  VN. 
The  element  is  prepared  by  heating  the  dichloride  in  a  stream  of  perfectly 
pure  hydrogen  — 


Vanadium  is  unacted  upon  by  air  at  ordinary  temperatures,  but  when 
heated  burns  brilliantly  to  the  pentoxide. 

Columbium  and  tantalum  are  found  associated  together  in  the  rare  mineral 
tantalite  or  columbite.  The  first  to  be  discovered  was  tantalum  (Hatchett, 
1801),  and  was  originally  named  columbium  ;  and  the  name  niobium  (from 
Niobe,  the  daughter  of  Tantalus)  was  given  to  the  allied  element  by  Rose 
(1846).  More  recently,  however,  the  name  of  this  second  element  has  been 
changed  to  columbium,  although  it  may  still  be  met  with  under  its  original 
name  of  niobium.  Columbium  is  obtained  by  heating  the  trichloride,  CbCl3, 
in  a  stream  of  hydrogen. 

Vanadium  forms  five  oxides,  corresponding  to  the  oxides  of  nitrogen,  while 
three  oxides  of  columbium  and  two  of  tantalum  are  known  :  — 


V20  ;   V20.,(or  VO) 
CbO 


V,,03  ;   V204(or  VO2)  ;   V2Oa 


CbOo 

Ta02         ;   Ta205. 

The  pentoxides  are  obtained  when  the  metals  are  burned  in  air  or  oxygen. 
They  give  rise  respectively  to  vanadates,  columbates,  and  tantalates,  cor- 
responding to  nitrates  and  metaphosphates,  thus — 

Sodium  nitrate,  NaNO3.  Sodium  metacolumbate,  NaCbO3. 

Sodium  metaphosphate,  NaPO3.        Sodium  metatantalate,  NaTaO3. 

Sodium  metavanadate,  NaVO3. 

655 


656  Inorganic  Chemistry 

The  closer  relation  of  these  elements  to  phosphorus  than  to  nitrogen  is  seen 
in  the  formation  of  salts  derived  from  ortho-  and  pyro-acids,  corresponding 
to  orthophosphates  and  pyrophosphates.  The  naturally  occurring  vanadium 
compounds  above  mentioned  are  vanadates  derived  from  the  hypothetical 
orthovanadic  acid,  H3VO4.  Both  metavanadic  acid,  HVO3,  and  pyrovanadic 
acid,  H4V2O7,  have  been  obtained.  Unlike  the  phosphorus  compounds,  the 
metavanadates  are  the  most  stable  of  the  three  classes  of  salts,  and  the 
orthovanadates  the  least  stable.  The  most  important  of  these  salts  is  the 
ammonium  metavanadate,  NH4VO3,  which  is  prepared  by  dissolving  the 
pentoxide  in  ammonia.  This  salt  is  insoluble  in  ammonium  chloride,  and 
use  is  made  of  this  property  in  the  preparation  of  vanadium  compounds 
from  the  mineral  mottramite.  When  ammonium  metavanadate  is  ignited, 
vanadium  pentoxide  is  obtained — 

2N  H4  V03 = V205  +  2N  H3  +  H2O. 

Vanadium  acts  also  as  a  feeble  base.  Thus,  when  the  tetroxide,  or  hypo- 
vanadic  oxide,  is  dissolved  in  sulphuric  acid,  hypovanadic  sulphate,  V2O2\SO4)2, 
is  formed.  The  solution  of  this  salt  possesses  a  rich  blue  colour. 

Vanadium  forms  three  chlorides,  having  the  composition — 

VCLj  (or  V2C14)  ;  VC13  (or  V2C16) ;  VC14. 

Columbinm  gives  a  trichloride,  CbCl3,  and  pentachloride,  CbCl5,  while  only 
the  pentachloride  of  tantalum  is  known,  TaCl5. 

Vanadium  forms  a  number  of  compounds  with  oxygen  and  chlorine.  Thus, 
when  vanadium  tetrachloride  is  acted  upon  by  water,  it  yields  hypovanadic 
chloride,  V2O4C12,  which  dissolves  in  the  water,  giving  a  blue  solution. 

Vanadium  oxychloride,  or  vanadyl  trichloride,  VOC13,  corresponds  to  phos- 
phorus oxychloride,  POC13.  From  vanadyl  trichloride,  by  treatment  with 
zinc,  vanadyl  dichloride  is  obtained,  VOC12,  and  by  the  action  of  hydrogen  at 
a  high  temperature  upon  this,  both  vanadyl  moncchloride,  VOC1,  and  divanadyl- 
monochloride,  V3O2C1,  are  formed. 


CHAPTER   XI 
ELEMENTS  OF  GROUP  VI.  (FAMILY  A.) 

Chromium,  Cr         .         .          .52.1  Tungsten,  W  184 

Molybdenum,  Mo    .  .96  Uranium,  U  239.5 

CHROMIUM. 

Symbol,  Cr.     Atomic  weight  =  52.  i. 

Occurrence. — Chromium  does  not  occur  in  nature  in  the  un- 
combined  state.  In  combination  with  oxygen  and  associated 
with  iron  it  is  met  with  in  considerable  quantities  in  the  mineral 
chrome  iron  ore,  or  chromite^  Cr2O3,FeO.  This  ore  is  the  chief 
source  of  chromium  compounds.  Other  natural  compounds  are 
crocoisite^  PbCrO4,  and  chrome- ochre,  Cr2O3.  Traces  of  chromium 
are  present  in  various  minerals,,  such  as  the  emerald  and  green 
serpentine,  and  impart  to  them  their  green  colour. 

Modes  Of  Formation. — Until  quite  recently  metallic  chromium 
was  a  mere  chemical  curiosity.  It  may  be  obtained  by  the  re- 
duction of  the  oxide,  Cr2O3,  by  means  of  carbon  at  the  high  tempera- 
ture of  the  electric  furnace.  The  metal  so  produced,  however, 
always  contains  carbon. 

It  is  now  produced  on  a  manufacturing  scale,  by  reducing  the 
oxide  by  means  of  metallic  aluminium.  The  powdered  oxide 
mixed  with  the  requisite  quantity  of  powdered  aluminium  is  placed 
in  a  refractory  crucible,  and  the  mixture  ignited  by  means  of  a  fuse. 

The  ignition  temperature  of  this  mixture  being  very  high,  the 
most  suitable  fuse  for  the  purpose  consists  of  a  mixture  of  barium 
peroxide  and  powdered  aluminium.  A  small  quantity  of  this 
mixture  is  placed  in  a  depression  made  in  the  surface  of  the  charge 
in  the  crucible,  and  a  piece  of  magnesium  ribbon  inserted  into  it. 
When  the  magnesium  is  ignited  it  immediately  fires  the  fuse,  which 
in  its  turn  communicates  its  combustion  to  the  charge.  The  con- 
tents of  the  crucible  undergo  rapid  vivid  combustion,  and  the  tem- 
perature of  the  entire  mass  rises  sufficiently  high  to  melt  the  reduced 
chromium. 

657 


658  Inorganic  Chemistry 

Properties. — Chromium  is  a  hard,  steel-grey  metal,  which  is 
not  oxidised  in  dry  air.  Its  melting-point  is  about  2000°,  being 
somewhat  higher  than  that  of  platinum.  The  metal  has  no  mag- 
netic properties.  It  dissolves  in  dilute  hydrochloric  and  sulphuric 
acids,  with  evolution  of  hydrogen.  When  placed  in  nitric  acid 
chromium  assumes  the  so-called  passive  condition,  and  while  in 
this  state  it  is  unacted  upon  by  the  acids  which  dissolve  it  under 
normal  conditions. 

Metallic  chromium  when  added  to  steel  imparts  to  the  latter  a 
high  degree  of  hardness  and  tenacity,  and  it  is  now  largely  em- 
ployed in  the  production  of  these  "chrome  steels,"  which  contain 
from  0.4  up  to  as  much  as  2  or  3  per  cent,  of  chromium. 

Oxides  of  Chromium. — Two  oxides  of  chromium  are  definitely 
known,  namely — 

Chromium  sesquioxide  (chromic  oxide)         .         .     Cr2O3. 
Chromium  trioxide  (chromic  anhydride]       .         .     CrO3. 

The  first  is  a  basic,  and  the  second  an  acidic  oxide.  Besides 
these  two  compounds,  a  hydrated  oxide,  derived  from  the  unknown 
chromous  oxide,  also  exists,  having  the  composition  CrO,H2O,  or 
Cr(HO)2.  It  is  obtained  as  a  yellowish  precipitate  by  adding  potas- 
sium hydroxide  to  a  solution  of  chromium  dichloride  (chromous 
chloride),  with  the  exclusion  of  air.  It  rapidly  absorbs  oxygen, 
turning  dark  brown.  When  heated  out  of  contact  with  air  it  is 
converted  into  the  sesquioxide,  with  evolution  of  hydrogen — 

2CrO,H2O  -  Cr2O3  +  H2O  +  H2. 

Other  compounds  of  chromium  and  oxygen  are  described,  whose  composi- 
tion, however,  is  not  definitely  established ;  thus,  the  product  obtained  as  a 
brown  powder,  either  by  the  partial  reduction  of  the  trioxide  or  the  oxidation 
of  the  sesquioxide,  is  regarded  by  some  chemists  as  chromium  dioxide,  CrO2, 
and  by  others  as  chromium  chromate,  Cr.2O3,CrO3.  It  is  readily  obtained  by 
passing  nitric  oxide  into  a  solution  of  potassium  dichromate. 

Chromium  Sesquioxide,  Cr2O3,  is  obtained  as  a  grey-green 
powder,  when  either  the  hydroxide,  or  the  trioxide,  or  ammonium 
dichromate  is  ignited  (see  page  230). 

When  the  vapour  of  chromyl  dichloride,  CrO2Cl2,  is  passed 
through  a  red-hot  tube,  chromic  oxide  is  deposited  in  the  form  of 
dark-green  hexagonal  crystals.  Chromic  oxide  which  has  been 
strongly  ignited  is  nearly  insoluble  in  acids.  It  is  used  under  the 
name  of  chrome  green  as  a  pigment,  and  for  giving  a  green  colour 
to  glass. 

Chromic   Hydroxides. — Chromic   oxide   yields   a    number   of 


Chromium   Trioxide  659 

hydrated  compounds.  When  ammonia  is  added  to  a  solution  of 
chromic  chloride,  or  other  chromic  salt,  free  from  alkali,  a  light 
blue  compound  is  precipitated,  which,  when  dried  over  sulphuric 
acid,  has  the  composition  Cr(HO)3,2H2O  (or  Cr2O3,7H2O).  When 
this  is  dried  in  vacuo  it  loses  water,  and  becomes  2[Cr(HO)3],H2O 
(or  Cr2O3,4H2O)  ;  and  on  being  heated  at  200°,  it  again  parts  with 
water,  and  has  the  composition  CrO(HO)  (or  Cr2O3,H.,O). 

When  potassium  dichromate  and  boric  acid  are  heated  to  dull 
redness,  and  the  mass  treated  with  water,  a  rich  green  residue 
is  obtained,  having  the  composition  Cr2O(HO)4  (or  Cr.,O3,2H2O). 
This  compound,  known  as  Guignefs  green,  is  employed  as  a 
pigment. 

The  first  two  of  these  compounds,  which  may  be  looked  upon 
as  consisting  of  the  hydroxide  Cr(HO)3  in  a  hydrated  condition, 
are  readily  soluble  in  acids,  yielding  the  chromic  salts. 

Chromium  Trioxide  {chromic  anhydride}  CrO3.  —  WThen  strong 
sulphuric  acid  is  added  to  a  cold  saturated  solution  of  potassium 
dichromate,  the  trioxide  separates  out  in  long,  red,  needle-shaped 
crystals  — 

K2Cr2O7  +  H2SO4=K2SO4+H2O  +  2CrO3. 

The  liquid  is  decanted  from  the  crystals,  which  are  drained 
upon  porous  tiles,  and  the  adhering  sulphuric  acid  and  potassium 
sulphate  washed  away  by  strong  nitric  acid.  The  crystals  are 
finally  heated  upon  a  sand-bath,  whereby  the  nitric  acid  is 
evaporated. 

Chromium  trioxide  dissolves  in  water  to  the  extent  of  62  parts 
in  100  parts  of  water  at  26°.  It  melts  at  a  temperature  about  192°. 
At  250°  it  begins  to  give  off  oxygen,  and  is  ultimately  converted 
into  the  sesquioxide  — 


Chromium  trioxide  is  a  powerful  oxidising  agent,  and  in  contact 
with  most  organic  substances  it  is  reduced.  In  the  preparation 
of  the  compound,  therefore,  the  liquid  cannot  be  filtered  through 
paper  in  the  usual  way.  Warm  alcohol  dropped  upon  the  trioxide 
at  once  takes  fire,  while  in  a  more  diluted  condition  it  is  oxidised 
to  acetic  acid  ;  and  the  reduction  of  the  chromium  trioxide  is  made 
evident  by  the  change  of  colour  of  the  liquid,  from  red  or  yellow 
to  olive  green. 

Gaseous  ammonia  reduces  the  trioxide  to  the  sesquioxide,  with 
formation  of  water  and  nitrogen  — 


660  Inorganic  Chemistry 

the  reaction  being  accompanied  with  the  evolution  of  so  much  heat 
that  the  chromic  oxide  produced  becomes  incandescent. 

When  hydrogen  peroxide  is  added  to  a  dilute  solution  of 
chromium  trioxide,  or  to  a  dilute  solution  of  potassium  dichromate, 
acidified  with  sulphuric  acid,  a  deep  indigo-blue  solution  is  ob- 
tained. This  blue  compound  is  believed  to  contain  perchromic 
acid,  but  its  composition  has  not  been  definitely  established.  It 
may  be  regarded  as  a  compound  of  chromium  trioxide,  CrO3,  or  of 
perchromic  acid,  HCrO4,  with  hydrogen  peroxide,  H2O2,  in  unde- 
termined proportions. 

In  aqueous  solution  the  blue  colour  quickly  disappears,  oxygen 
being  eliminated.  The  compound  is  soluble  in  ether  ;  and  there- 
fore, when  the  aqueous  solution  is  shaken  up  with  that  liquid,  a 
deep  blue  ethereal  solution  rises  to  the  top.  In  this  solution  the 
compound  is  more  stable,  but  when  evaporated  it  evolves  oxygen, 
leaving  chromium  trioxide.  It  is  decomposed  by  alkalies,  forming 
alkaline  chromates  with  evolution  of  oxygen.  The  formation  of 
this  compound  constitutes  a  delicate  test  for  either  chromium 
trioxide  or  hydrogen  peroxide  (see  Hydrogen  Peroxide,  page  227). 

Chromous  Compounds. — These  correspond  to  chromous hydrate,  Cr(HO)2, 
in  which  the  chromium  functions  as  a  divalent  element.  Comparatively  few 
of  these  salts  are  known. 

Chromous  Chloride,  CrCI2,  is  formed  when  the  metal  dissolves  in  hydro- 
chloric acid.  It  is  prepared  in  the  anhydrous  state  by  gently  heating  chromic 
chloride  in  a  current  of  pure  hydrogen.  It  is  a  white  crystalline  compound, 
soluble  in  water  to  a  blue  solution,  which  rapidly  absorbs  oxygen. 

Chromous  Sulphate,  CrSO^7H2O,  is  obtained  by  dissolving  chromous 
acetate  in  dilute  sulphuric  acid.  It  is  deposited  from  the  solution  in  blue 
crystals,  isomorphous  with  ferrous  sulphate,  FeSO4,7HoO. 

Chromic  Compounds. — These  are  derived  from  chromic  oxide, 
the  oxide  acting  as  a  base. 

Chromic  Chloride,  CrCl3,  or  Cr2Cl6,  is  prepared  by  strongly 
heating  a  mixture  of  chromic  oxide,  Cr2O3,  and  carbon  in  a  stream 
of  dry  chlorine.  The  chromic  chloride  sublimes  in  the  form  of 
scales,  having  a  reddish-pink  colour.  The  molecular  weight  of 
chromic  chloride  is  158.45,  showing  that  in  the  vaporous  state  its 
molecules  have  the  formula  CrCl3. 

It  is  nearly  insoluble  in  water,  but  readily  dissolves  in  water 
containing  minute  traces  of  chromous  chloride,  forming  a  green 
solution.  The  same  solution  is  obtained  by  dissolving  hydrated 
chromic  hydroxide,  Cr(HO)3,2H2O,  in  hydrochloric  acid,  and  if 
this  solution  be  slowly  evaporated,  very  soluble  green  crystals 
separate  out,  having  the  composition  CrCl3,6H2O.  If  strongly 
heated  in  the  air,  this  compound  gives  off  water  and  hydrochloric 


Chrome  Alum  66  1 

acid,  leaving  chromic  oxide,  Cr2O3  ;  but  when  heated  to  250°,  in 
either  gaseous  hydrochloric  acid  or  chlorine,  it  is  converted  into 
the  pink  anhydrous  chromic  chloride,  which  redissolves  in  water 
to  the  green  solution.  If  heated  strongly  and  sublimed,  the  com- 
pound obtained  is  nearly  insoluble  in  water. 

Chromic  Sulphate,  Cr2(SO4)3,  is  obtained  by  dissolving  chro- 
mium hydroxide  in  concentrated  sulphuric  acid,  when  a  green 
solution  is  formed,  which  on  standing  changes  to  blue,  and  slowly 
deposits  violet-blue  crystals.  The  salt  may  be  purified  by  dis- 
solving in  cold  water  and  precipitating  with  alcohol.  If  insufficient 
alcohol  be  added  to  cause  immediate  precipitation,  the  salt  slowly 
deposits  from  the  dilute  spirit  in  blue  octahedrons,  belonging  to 
the  regular  system. 

A  cold  aqueous  solution,  which  has  a  violet  colour,  becomes 
green  when  boiled. 

Chromic  sulphate  forms  double  salts  with  the  sulphates  of  the 
alkalies,  which  belong  to  the  alums. 

Potassium  Chromium  Alum  (chrome  alum],  K2SO4,Cr2(SO4)3, 
24H2O.  —  This  double  sulphate  is  formed  when  solutions  of  potas- 
sium and  chromium  sulphates  are  mixed  together  in  molecular 
proportions.  It  is  most  conveniently  prepared  by  the  addition  of 
the  requisite  amount  of  sulphuric  acid  to  an  aqueous  solution  of 
potassium  dichromate,  and  reducing  the  chromic  oxide  by  passing 
sulphur  dioxide  through  the  liquid  — 


(1)  K2Cr2O7  +  H2SO4 

(2)  2CrO3  +  3SO2    =Cr2(S04)3. 

The  resulting  solution,  containing  the  two  sulphates  in  mole- 
cular proportions,  deposits  crystals  of  the  double  sulphate,  in  the 
form  of  dark  plum-coloured  octahedrons  (Fig.  147,  B,  page  619), 
which  appear  red  by  transmitted  light. 

Chrome  alum  dissolves  in  water,  yielding  a  plum-coloured  solu- 
tion, which  on  boiling  turns  green,  but  on  long  standing  returns  to 
its  original  colour. 

Sodium  chromium  alum  is  more  soluble,  and  ammonium  chro- 
mium alum  is  less  soluble,  than  the  potassium  salt. 

Chromites.  —  Chromic  oxide  acts  also  as  a  weak  acid,  and  combines  with 
other  oxides,  forming  compounds  resembling  the  aluminates.  When  potas- 
sium hydroxide  is  added  to  a  solution  of  a  chromic  salt,  the  green  hydrated 
oxide  which  is  precipitated  contains  alkali  which  cannot  be  removed  by  hot 


662  Inorganic  Chemistry 

water ;  this  is  present  in  the  form  of  potassium  chromite.  The  best  known 
chromites  are  zinc  chromite,  Cr2O3,ZnO  ;  manganous  chromite,  Cr2O3,MnO, 
and  ferrous  chromite,  Cr2O3,FeO;  the  latter  occurs  naturally  as  chrome 
iron  ore. 

Chromates. — When  chromium  trioxide  is  dissolved  in  water, 
the  solution  is  believed  to  contain  chromic  acid,  H2CrO4,  or 
dichromic  acid,  H2Cr2O7  ;  when  the  solution  is  evaporated,  how- 
ever, the  trioxide  alone  is  left.  Red  crystals  have  been  obtained, 
by  cooling  a  hot  saturated  solution  of  the  trioxide,  which  are 
believed  to  be  chromic  acid. 

Potassium  Chromate,  K2CrO4,  is  prepared  by  adding  potas- 
sium hydroxide  to  a  solution  of  the  dichromate — 

K2CroO7  +  2KHO  =  2K2CrO4+H2O. 

On  evaporation,  the  yellow  chromate  of  potash  separates  out  in 
rhombic  crystals,  isomorphous  with  potassium  sulphate.  It  is 
soluble  in  water  at  the  ordinary  temperature  to  the  extent  of  60 
parts  in  100  parts  of  water,  forming  a  yellow  solution  having  an 
alkaline  reaction. 

Potassium  Bichromate,  K2Cr2O7,  is  manufactured  from  chrome 
iron  ore  by  roasting  the  finely  crushed  ore  with  potassium  car- 
bonate and  lime  in  a  reverberatory  furnace  ;  the  mass  being 
frequently  raked  over  to  expose  fresh  portions  to  the  oxidising 
action  of  the  flames.  In  this  way  a  mixture  of  calcium  and  potas- 
sium chromates  is  produced — 

2Cr2O3,  FeO  +  3KrCO3  +  CaO  +  7O = CaCrO4 + 3K.2CrO4  +  Fe2O3  +  3CO.2. 

The  yellow  mass,  when  cold,  is  broken  up  and  lixiviated  with  a 
hot  solution  of  potassium  sulphate,  which,  by  double  decomposition 
with  the  calcium  chromate,  forms  potassium  chromate  and  precipi- 
tates calcium  sulphate.  The  solution  after  settling  is  treated  with 
the  requisite  quantity  of  sulphuric  acid  to  convert  the  chromate 
into  the  dichromate,  thus — 

2K2CrO4+  H2SO4  =  K2SO4  +  H2O  +  K2Cr2Or. 

The  dichromate  being  much  less  soluble  than  the  normal  chro- 
mate, a  large  proportion  of  it  at  once  deposits  as  the  solution  cools  ; 
and  the  mother-liquor  containing  potassium  sulphate  is  used  again 
to  lixiviate  a  fresh  quantity  of  the  roasted  mixture. 

Potassium  dichromate  forms  large  red  prisms  or  tables,  belong- 


Chromyl  Chloride  663 

ing  to  the  asymmetric  (triclinic)  system.  It  is  soluble  in  water  at 
the  ordinary  temperature  to  the  extent  of  10  parts  in  100  parts  of 
water,  yielding  an  acid  solution,  which  is  extremely  poisonous. 
When  a  film  of  gelatine  is  impregnated  with  potassium  dichromate 
and  exposed  to  light,  a  reduction  of  the  chromium  to  chromic 
oxide  takes  place,  and  at  the  same  time  the  gelatine  is  rendered 
insoluble.  This  property  is  utilised  in  photographic  processes.* 

Potassium  dichromate  is  also  known  under  the  misnomer  bichromate  of 
potash,  which  would  suggest  that  the  salt  was  in  reality  hydrogen  potassium 
chromate,  corresponding  to  bisulphate  of  potash,  HKSO4.  Such  a  chromium 
compound  does  not  exist.  The  dichromates  correspond  to  the  disulphates  (or 
pyrosulphates),  see  page  435. 

Potassium  Tricliromate,  K2Cr3O10  (or  K2CrO4l2CrO3),  and  Potassium 
Tetrachromate,  K2Cr4O13(or  KaCrO4,3CrO3),  are  also  known. 

Lead  Chromate,  PbCrO4,  is  found  as  the  mineral  crocoisite. 
It  is  produced  by  precipitation  from  a  lead  salt,  with  either 
potassium  chromate  or  dichromate.  It  forms  a  bright  yellow 
powder,  known  as  chrome-yellow ',  and  is  employed  as  a  pigment. 
It  melts  without  decomposition,  and  resolidifies  on  cooling  to  a 
brown  crystalline  solid.  At  higher  temperatures  it  gives  off 
oxygen,  and  is  converted  into  chromic  oxide  and  a  basic  lead 
chromate.  When  heated  with  organic  compounds,  the  latter  are 
completely  oxidised ;  lead  chromate  is  therefore  employed  in 
organic  analyses. 

When  lead  chromate  is  digested  with  sodium  hydroxide,  or  with 
normal  potassium  chromate,  a  basic  lead  chromate  is  obtained 
as  a  rich  red  powder — 

2PbCrO4  +  2NaHO  =  Na2CrO4  +  H2O  +  Pb2CrO5. 

This  compound  is  known  as  chrome-red. 

Chromyl  Chloride,  CrO2Cl2.— This  compound  is  prepared  by 
distilling  a  mixture  of  potassium  dichromate  and  sodium  chloride 
with  strong  sulphuric  acid.  Chromyl  chloride  is  a  deep  red,  mobile, 
strongly  fuming  liquid.  It  is  decomposed  by  water  into  hydro- 
chloric acid  and  chromium  trioxide,  and  acts  as  a  powerful 
oxidising  substance.  When  dropped  upon  phosphorus  it  explodes. 
When  heated  in  sealed  tubes  it  is  converted  into  trichromyl 
chloride  with  loss  of  chlorine,  (CrO2)3Cl2. 

Chromyl   chloride   may   be    regarded    as   being   derived   from 

*  Abney,  "  Treatise  on  Photography." 


664  Inorganic  Chemistry 

chromic  acid,  CrO2(HO)2,  by  the  complete  substitution  of 
(HO)  by  Cl.  The  intermediate  compound,  chloro-chromic  acid, 
CrO2(HO)Cl,  is  unknown,  although  its  salts  have  been  pre- 
pared ;  thus,  by  the  gentle  action  of  hydrochloric  acid  upon 
potassium  dichromate,  potassium  chloro-chromate  is  obtained  as 
a  red  crystalline  salt — 


:  =  2CrO2(KO)Cl  +  H2O. 

Molybdenum,  Mo  =  96;  Tungsten,  W  =  i84;   Uranium,  U=2^8.^. 

These  three  somewhat  rare  elements  are  closely  related  to  chromium. 

Molybdenum  occurs  in  the  mineral  molybdenite,  MoS.2  (which  strongly  re- 
sembles graphite  in  appearance),  and  more  rarely  as  molybdenum  ochre,  MoO3, 
and  wulfenite,  PbMoO4. 

Tungsten  is  found  chiefly  in  wolfram,  2FeWO4,3MnWO4  (occurring  in 
the  Cornish  tin  mines) ;  more  rarely  as  scheelinite,  PbWO4,  and  wolfram 
ochre,  WO3. 

Uranium  occurs  as  an  oxide,  UO.2,2UO3,  in  pitchblende  (a  considerable 
quantity  of  which,  associated  with  other  uranium  compounds,  has  recently 
been  discovered  at  St.  Stephens,  Cornwall). 

Molybdenum  is  obtained  by  the  action  of  hydrogen  upon  the  heated  oxide 
or  chloride ;  uranium,  by  the  action  of  sodium  upon  the  chloride ;  while 
tungsten  has  been  obtained  by  both  methods.  In  their  specific  gravities, 
tungsten  and  uranium  exhibit  a  marked  difference  from  chromium  and 
molybdenum;  thus,  Cr,  sp.  gr.  =6 ;  Mo,  sp.  gr.  =8.6;  while  W,  sp.  gr.  = 
19.  i  ;  U,  sp.  gr.  =  18.7. 

Molybdenum' and  uranium  form  a  large  number  of  oxides,  some  of  which 
are  regarded  as  definite  oxides,  while  others  are  looked  upon  as  combinations 
of  two  oxides.  Only  two  oxides  of  tungsten  are  known.  The  composition  of 
these  compounds  is  as  follows — 

MoO 

Mo.203 

Mo02  W02  J02    U2o5=U02UO3 

IVlOC/Q  \\   O«J  LJ  O3  T  T      /*\  TT/-\        f\T  T^-N 

_  3  _3  U04    U30S=U02,2UO3. 

The  trioxide  of  each  metal  is  an  acid  oxide ;  uranium  trioxide,  however,  is 
both  acidic  and  basic.  .  They  are  insoluble  in  water,  but  by  the  action  of 
alkalies  they  yield  molybdates,  tungstates,  and  uranates.  Molybdates  and 
tungstates,  derived  from  the  acids  H2MoO4,2H2O  and  H2WO4,2H2O  (corre- 
sponding to  chromic  acid),  are  known.  And  all  three  oxides  yield  salts 
corresponding  to  potassium  dichromate,  thus — 

Sodium  Dimolybdate.  Sodium  Ditungstate.  Sodium  Diuranate. 

Na2Mo.2O7  Na2W2O7  Na.7U0O7. 


Molybdic  and  tungstic  acids  also  form  numerous  polymolybdates  and  poly- 


Molybdenum,  Tungsten ,   Uranium  665 

tungstates,  by  the  absorption  of  varying  quantities  of  the  trioxide  into  the 
molecule  of  the  normal  salt  (see  Chromates,  page  663).  And  in  the  case  of 
tungsten,  the  compound  metatungstic  acid,  H2W4O13,7H2O,  is  known. 

Uranium  dioxide  and  trioxide  are  both  basic  oxides,  the  former  yielding  the 
unstable  uranous  salts,  such  as  uranous  sulphate,  U(SO4)2;  and  the  latter 
producing  the  uranyl  salts,  of  which  the  sulphate,  (UO2)SO4,  and  nitrate, 
(UOo)(NO3)2,  are  well  known. 

Uranium  peroxide,  UO4,  is  an  acid  oxide  which  yields  per-uranates. 

Both  molybdic  and  tungstic  acids  form  complex  compounds  with  phos- 
phoric acid,  known  as  phospho-molybdic  and  phospho-tungstic  acids  :  thus, 
when  a  nitric  acid  solution  of  ammonium  molybdate  (NH4)2MoO4,  is  added 
in  excess  to  a  solution  of  orthophosphoric  acid  or  an  orthophosphate,  a 
canary- yellow  crystalline  precipitate  of  ammonium  phospho-molybdate, 
2(NH4)3PO4,22MoO3,12H2O,  is  obtained  (see  page  477).  It  is  soluble  in 
alkalies  and  in  excess  of  phosphoric  acid,  but  insoluble  in  dilute  mineral  acids. 
\Yhen  this  compound  is  dissolved  in  aqua-regia  the  solution  deposits  yellow 
crystals  of  phospho-molybdic  acid,  2H3PO4,22MoO3. 

Compounds  with  chlorine  having  the  following  composition  are  known — 

MoCl2  WC12 

MoCl3  or  MooClc 

MoCl4  WC14  UC14. 

MoCls  WC15  UClg. 


CHAPTER  XII 
GROUP  VII.  (FAMILY  A.} 

MANGANESE. 

Symbol,  Mn.     Atomic  weight  =  54. 93. 

Occurrence. — This  element  is  never  found  in  nature  in  the  free 
state.  It  is  widely  distributed  in  combination  with  oxygen,  as 
Pyrolusite,  MnO2 ;  braunite,  Mn2O3 ;  and  hausmannite^  Mn3O4. 
Also  as  a  hydrated  oxide  in  manganite,  Mn2O3,H2O.  It  is  met 
with  also  as  carbonate  in  manganese  spar,  MnCO3 ;  and  as  sul- 
phide in  manganese  blende,  MnS. 

Modes  Of  Formation.— Manganese  may  be  obtained  by  the 
reduction  of  the  oxide  by  means  of  carbon  at  a  very  high  tempera- 
ture, as  obtained  in  the  electric  furnace.  The  product,  however, 
contains  carbon.  In  a  purer  state  it  may  be  prepared  by  the  re- 
duction of  fused  anhydrous  manganous  chloride  by  means  of  metallic 
magnesium.  At  the  present  time,  however,  it  is  obtained  by  reduc- 
tion of  the  oxide  by  means  of  aluminium,  in  a  manner  precisely 
similar  to  that  employed  for  the  manufacture  of  chromium. 

Properties. — Manganese  is  a  hard,  brittle  metal,  the  colour  of 
which  is  grey  with  a  slight  tinge  of  red.  Its  melting-point  is  below 
that  of  chromium  but  higher  than  that  of  iron,  being  about  1900°. 
It  slowly  oxidises  on  exposure  to  moist  air,  and  is  readily  dissolved 
by  dilute  sulphuric,  hydrochloric,  or  even  acetic  acid,  with  evolution 
of  hydrogen.  The  chief  use  of  manganese  is  in  the  iron  industry 
(see  Iron). 

Oxides  Of  Manganese. — The  four  most  important  of  these  are — 
Manganous  oxide  (manganese  monoxide)    .         .     MnO. 
Red  manganese  oxide  ......     Mn3O4. 

Manganic  oxide  (manganese  sesquioxide)    .         .     Mn2O3. 
Manganese  dioxide     c          .....     MnO2. 

The  monoxide  and  sesquioxide  are  basic,  giving  rise  to  man- 
ganous and  manganic  salts  respectively.  The  oxide,  Mn3O4,  is 
also  basic,  but  yields  with  acids  both  manganous  and  manganic 

666       • 


Manganese  Dioxide  667 

salts.  Manganese  dioxide  or  peroxide,  MnO2,  gives  manganous 
salts  with  elimination  of  available  oxygen.  It  also  combines  with 
certain  more  basic  oxides,  forming  unstable  compounds  known  as 
inanganites. 

Manganese  trioxide,  MnO3,  and  hept-oxide,  Mn2O7,  have  also 
been  obtained.  They  are  both  acid  oxides,  giving  rise  respec- 
tively to  the  manganates  &\i&permanga?iates. 

Manganous  Oxide,  MnO,  is  obtained  by  heating  any  of  the 
higher  oxides  in  a  stream  of  hydrogen,  or  by  igniting  a  mixture  of 
manganous  chloride,  sodium  carbonate,  and  ammonium  chloride. 
It  is  a  light  green  powder,  which,  if  prepared  at  a  low  tempera- 
ture, oxidises  in  the  air.  When  perfectly  air-free  solutions  of 
potassium  hydroxide  and  a  manganous  salt  are  mixed,  with  exclu- 
sion of  air,  hydrated  manganous  oxide,  or  manganous  hydroxide, 
Mn(HO)2,  is  obtained  as  a  white  precipitate,  which  rapidly  oxidises 
on  exposure  to  air. 

Red  Manganese  Oxide  (mangano-manganic  oxide),  Mn3O4,  is 
the  most  stable  of  the  oxides  of  manganese,  being  formed  when 
both  the  higher  or  lower  oxides  are  strongly  heated.  Thus,  in  the 
preparation  of  oxygen  by  heating  the  dioxide,  this  compound 
remains  (page  184).  With  cold  sulphuric  acid  it  yields  a  mix- 
ture of  manganous  and  manganic  sulphates,  but  when  heated  with 
dilute  acid,  manganous  sulphate  and  dioxide  are  formed — 

Mn3O4  +  2H2S04  =  2MnSO4+MnO,  +  2H20. 

Manganic  Oxide  (manganese  sesquioxide),  Mn2O3,  occurs  native 
as  braunite,dc(\&  in  the  hydrated  condition  as  manganite,  Mn2O3,H2O. 
The  hydrated  oxide  is  formed  by  the  spontaneous  oxidation  of  man- 
ganous hydroxide,  and  when  gently  heated  it  yields  the  oxide. 
Both  the  oxide  and  the  hydrate,  on  treatment  with  warm  nitric 
acid,  yield  manganous  nitrate  and  manganese  dioxide. 

Manganese  Dioxide,  MnO2,  is  the  most  important  of  the  man- 
ganese ores.  It  may  be  obtained  by  the  cautious  ignition  of 
manganous  nitrate — 

Mn(NO3)2=2NO2  +  MnO.j. 

Manganese  dioxide  is  a  hard  black  solid  which  conducts  electri- 
city and  is  strongly  electro-negative  to  metals.  On  this  account 
it  is  employed  in  certain  forms  of  voltaic  battery.  When  heated 
it  loses  oxygen,  and  forms  first  the  sesquioxide  and  finally  Mn3O4. 


668  Inorganic  Chemistry 

Manganese  dioxide  dissolves  in  cold  concentrated  hydrochloric 
acid,  forming  a  dark-brown  solution  which  is  believed  to  contain 
the  compound  MnCl3.  On  warming  it  evolves  chlorine,  and  leaves 
manganous  chloride,  MnCl2. 

Manganites. — Manganese  dioxide  combines  with  certain  me- 
tallic oxides,  forming  unstable  compound  oxides.  Thus,  with  lime 
it  forms  CaO,MnO2 ;  CaO,2MnO2,  and  CaO,5MnO2.  These  com- 
pounds are  produced  in  the  V/eldon  recovery  process  (page  359). 

MANGANOUS  SALTS. 

Manganous  Chloride,  MnCl2,  is  the  only  chloride  of  this  metal 
that  has  been  isolated.  It  is  obtained  by  dissolving  any  of  the 
oxides  or  the  carbonate  in  hydrochloric  acid,  and  on  evaporation 
is  deposited  as  pink  crystals  of  MnCl2,4H2O.  The  anhydrous 
salt  is  prepared  by  heating  the  crystals  in  a  stream  of  hydro- 
chloric acid.  Manganese  chloride  forms  double  salts  with  chlo- 
rides of  the  alkalies,  the  ammonium  salt  MnCl2,2NH4Cl,H2O  being 
the  best  known. 

Manganous  Sulphate,  MnSCX,  is  prepared  by  strongly  heating 
a  pasty  mixture  of  the  dioxide  and  strong  sulphuric  acid.  The 
iron  present  is  thereby  converted  into  ferric  oxide,  and  on  treat- 
ing the  calcined  mass  with  water  manganous  sulphate  dissolves. 
The  solution  on  evaporation  deposits,  at  ordinary  temperatures, 
large  pink  crystals  of  MnSO4,5H2O  (isomorphous  with  copper 
sulphate).  Below  6°  rhombic  crystals  are  formed  (also  pink)  of  the 
composition  MnSO4,7H2O  (isomorphous  with  ferrous  sulphate). 

When  these  salts  are  heated  to  200°,  or  when  their  solutions  are 
boiled,  the  anhydrous  salt  is  formed.  With  sulphates  of  the 
alkalies,  manganous  sulphate  forms  double  salts,  as  potassium 
manganous  sulphate,  K2SO4,MnSO4,6H2O  ;  and  with  aluminium 
sulphate  it  yields  &  pseudo-alum  (see  page  620),  MnSO4,Al2(SO4)3, 
24H2O. 

MANGANIC  SALTS. 

Manganic  Chloride  is  obtained  as  a  dark-brown  solution 
when  the  dioxide  is  dissolved  in  cold  hydrochloric  acid.  It  has 
never  been  isolated,  and  is  believed  to  have  the  composition 
MnCLj. 

Manganic  Sulphate,  Mn2(SO4)3,  is  obtained  as  a  green  deli- 
quescent powder  by  the  action  of  sulphuric  acid  upon  the  pre- 


Permanganates  669 

cipitated  peroxide.  On  exposure  to  the  air  the  deliquesced  mass 
becomes  muddy,  by  the  precipitation  of  hydrated  manganic  oxide, 
thus  — 

=  3H2SO4+Mn2O3,H20. 


On  the  addition  of  potassium  sulphate  to  a  solution  of  manganic 
sulphate  in  dilute  sulphuric  acid,  potassium  manganese  alum  is 
obtained,  K2SO4,Mn2(SO4)3,24H2O,  which  deposits  in  violet  regular 
octahedra.  In  the  presence  of  much  water  the  salt  is  decomposed, 
and  deposits  the  hydrated  manganic  oxide. 

MANGANATES. 

These  salts  are  derived  from  the  hypothetical  manganic  acid, 
H2MnO4.  The  oxide  corresponding  to  this  acid  is  known,  viz., 
MnO3.  It  is  an  unstable  compound,  obtained  as  a  reddish  amor- 
phous mass,  by  adding  a  solution  of  potassium  permanganate  in 
sulphuric  acid  to  dry  sodium  carbonate. 

The  manganates  of  the  alkalies  are  obtained  by  fusing  manganese 
dioxide  with  potassium  or  sodium  hydroxide.  If  air  be  excluded 
the  following  reaction  takes  place  — 


In  the  presence  of  air  or  oxygen,  or  by  the  addition  of  potassium 
nitrate  or  chlorate,  more  of  the  manganese  is  converted  into  man- 
ganate.  The  fused  mass  has  a  dark-green  colour,  and  dissolves  in 
a  small  quantity  of  cold  water  to  a  deep  green  solution,  which  is 
only  stable  in  the  presence  of  free  alkali. 

When  a  solution  of  potassium  manganate  is  largely  diluted  or 
gently  warmed,  it  changes  from  green  to  pink,  owing  to  the  con- 
version of  the  manganate  into  permanganate,  thus  — 


The  same  change  takes  place  when  carbon  dioxide  is  passed 
through  the  solution. 

PERMANGANATES. 

These  salts  are  derived  from  permanganic  acid,  HMnO4.  When 
potassium  permanganate  is  cautiously  added  to  cold  strong  sul- 
phuric acid,  green  oily  drops  of  the  unstable  manganese  heptoxide 


6/O  Inorganic  Chemistry 

(or  permanganic  anhydride)  are  obtained,  Mn2O7.  This  compound 
dissolves  in  a  small  quantity  of  water  to  a  purple  solution,  which 
contains  the  unstable  acid,  Mn2O7,H2O,  or  H2Mn2O8  =  2HMnO4 
The  solution  evolves  oxygen  and  deposits  manganese  dioxide. 

Potassium  Permanganate,  KMnO4,  is  the  most  important  salt 
of  this  class.  It  is  prepared  by  fusing  the  dioxide  with  potassium 
hydroxide  and  potassium  chlorate,  dissolving  the  manganate  so 
obtained  in  water,  and  passing  carbon  dioxide  through  the  solu- 
tion. The  filtered  solution,  on  evaporation,  deposits  dark  purple 
rhombic  prisms,  which  appear  deep  red  by  transmitted  light. 
Potassium  permanganate  is  isomorphous  with  potassium  per- 
chlorate,  KCKX ;  it  dissolves  in  water,  forming  a  rich  purple 
solution.  When  boiled  with  strong  caustic  alkalies  it  loses  oxygen 
and  forms  the  green  potassium  manganate — 

2KMnO4  +  2KHO  =  2K2MnO4+H2O  +  O. 

It  readily  gives  up  oxygen  to  oxidisable  and  organic  compounds, 
and  on  this  account  is  used  both  as  a  laboratory  oxidising  agent 
and  as  a  disinfectant.  The  crude  sodium  salt  is  largely  employed, 
under  the  name  of  Condy*s  Disinfecting  Fluid,  for  this  purpose. 
When  solid  potassium  permanganate  is  heated  to  240°  it  evolves 
oxygen,  and  forms  potassium  manganate  and  manganese  dioxide — 

2KMnO4=  K2MnO4+  MnO2  +  O2. 


CHAPTER   XIII 

THE    TRANSITIONAL   ELEMENTS   OF   THE    FIRST 
LONG    PERIOD 

Iron,  Fe  =  ss.8s.         Cobalt,  €0  =  58.97.         Nickel,  Ni= 58. 7. 

THESE  three  elements  belonging  to  Group  VIII.  (see  classifica- 
tion, page  118)  stand  in  a  different  relation  to  each  other  than  the 
members  of  the  other  seven  groups. 

Iron,  cobalt,  and  nickel  belong  to  the  same  period,  being  the 
transitional  elements  falling  between  the  first  and  second  series  of 
the  first  long  period.  They  are  related,  on  the  one  hand,  through 
iron,  to  the  preceding  metals  manganese  and  chromium  (see  such 
compounds  as  ferrates,  manganates,  chromates] ;  while,  on  the  other 
hand,  through  nickel,  they  approach  the  metal  copper,  which  is  the 
next  following  in  the  period. 

Iron,  cobalt,  and  nickel  are  closely  related  elements  ;  in  nature 
they  are  usually  associated  together.  They  are  all  attracted  by 
the  magnet,  and  are  nearly  white,  hard,  and  difficultly  fusible 
metals.  In  their  chemical  habits,  however,  they  exhibit  a  gradual 
transition  in  their  properties.  Thus,  iron  forms  two  basic  oxides, 
yielding  two  series  of  stable  salts,  viz.,  ferrous  and  ferric.  Cobalt 
also  has  two  basic  oxides,  but  the  basicity  of  the  sesquioxide  is 
very  feeble,  and  cobaltzV  salts  (except  double  salts)  are  unstable, 
and  are  only  known  in  solution.  Nickel  only  forms  one  basic 
oxide,  and  yields  only  one  series  of  salts  corresponding  to  the 
ferrous  salts,  the  sesquioxide  of  nickel  behaving  with  acids  as  a 
peroxide. 

IRON. 

Symbol,  Fe.      Atomic  weight=55.85. 

Occurrence. — Iron  is  one  of  the  most  abundant  and  widely 
distributed  elements.  It  occurs  in  the  uncombined  state  in  small 

particles  disseminated  through  certain  basalts,  and  also  in  meteoric 

671 


672  Inorganic  Chemistry 

iron,  where  it  is  usually  associated  with  nickel,  cobalt,  and  copper. 
Masses  of  iron  have  also  been  found  which  have  been  formed  by 
the  reduction  of  iron  ores,  owing  to  the  firing  of  coal  pits  :  such 
iron  is  known  as  natural  steeL 

The  chief  ores  of  iron  are  red  h<zmatite  and  specular  iron  ore, 
Fe2O3;  brown  htzmatite,  2Fe2O3,3H2O  ;  magnetic  iron  ore  (load- 
stone), Fe3O4  ;  spathic  iron  ore,  FeCO3  ;  clay  iron  stone  consists 
of  spathose  iron  mixed  with  clay  ;  and  blackband  is  clay  iron 
stone  containing  from  20  to  25  per  cent,  of  coal. 

Iron  is  also  found  in  combination  with  sulphur,  as  iron  pyrites, 
FeS2,  and  with  iron  and  copper  in  copper  pyrites,  Cu2S,Fe2S3,  but 
these  compounds  are  not  employed  in  the  metallurgy  of  iron. 

Modes  of  Formation.  —  Iron  is  readily  reduced  from  its  com- 
pounds. Thus,  if  ferric  oxide  or  oxalate  be  gently  heated  in  a 
stream  of  hydrogen,  the  metal  is  obtained  as  a  black  powder, 
which  spontaneously  oxidises  with  incandescence  when  brought 
into  the  air.  On  the  industrial  scale  the  reduction  is  effected  by 
means  of  coke  and  limestone.  The  ore  is  first  calcined,  whereby 
water  and  carbon  dioxide  are  expelled,  and  any  sulphides  present 
are  oxidised,  with  the  expulsion  of  sulphur  dioxide.  By  this  pro- 
cess also  the  ore  is  rendered  more  porous.  The  calcined  ore  is 
then  smelted  in  a  blast-furnace,  with  limestone  and  coke.  Fig.  1  53 
shows  in  section  a  modern  blast-furnace.  The  charge  is  admitted 
at  the  top  by  means  of  the  cup  and  cone  arrangement,  which  closes 
the  furnace,  and  a  powerful  hot-blast  is  forced  through  tuyeres 
placed  round  the  base  of  the  furnace.  The  furnace  gases  are  led  off 
by  the  side  pipe  at  the  top,  and  are  utilised  for  heating  the  blast. 

The  chemical  reactions  which  take  place  in  a  blast-furnace  are 
many  and  complex,  and  differ  in  different  parts  of  the  furnace. 
In  the  main,  the  following  are  the  changes  which  occur.  The 
atmospheric  oxygen  of  the  hot-blast,  on  coming  in  contact  with 
the  carbon,  forms  carbon  monoxide  (at  the  high  temperature 
carbon  dioxide  is  probably  not  first  formed).  As  the  charges  of 
ore  gradually  work  their  way  down  the  furnace,  they  soon  arrive 
at  a  point  where  the  ferric  oxide  begins  to  be  reduced  by  the 
heated  carbon  monoxide,  first  to  ferrous  oxide,  and  then  to  a 
spongy  or  porous  mass  of  metallic  iron.  The  region  where  this 
takes  place  is  termed  the  zone  of  reduction  — 


In  the  early  stages  of  its  descent  through  the  furnace,  the  lime 


Iron 


673 


stone  is  converted  into  carbon  dioxide  and  lime.  The  reduced 
spongy  metal,  as  it  passes  down  through  the  hotter  regions  of  the 
furnace,  begins  to  take  up  carbon.  It  is  probable  that  carbon 


15-0 > 


FIG.  153. 

monoxide  first  combines  with  the  reduced  iron,  forming  iron 
carbonyl  (see  page  300),  which  at  a  higher  temperature  is  decom- 
posed, with  the  precipitation  of  finely  divided  carbon  within  the 
pores  of  the  mass.  More  and  more  carbon  is  taken  up  by  the  iron 

2  U 


674  Inorganic  Chemistry 

as  it  descends,  until  it  passes  from  a  pasty  condition  to  a  state  of 
complete  fusion,  when  it  collects  upon  the  bottom,  or  hearth,  of  the 
furnace.  In  passing  through  the  hottest  regions  the  lime  combines 
with  the  siliceous  materials  originally  present  in  the  ore  to  form  a 
fusible  slag,  beneath  which  the  molten  iron  collects.  Other  re- 
actions which  go  on  in  various  regions  of  the  furnace  are  the  reduc- 
tion of  sulphur  compounds,  and  of  phosphates  and  silicates,  with 
the  absorption  into  the  iron  of  a  certain  amount  of  sulphur,  phos- 
phorus, and  silicon.  The  precise  nature  of  the  changes  suffered 
by  the  gases  in  the  various  regions  of  the  furnace  is  still  obscure. 
The  cyanogen  formed  by  the  direct  union  of  atmospheric  nitrogen 
with  carbon,  and  also  the  hydrocarbons  present,  doubtless  undergo 
a  chemical  change  in  contact  with  the  heated  iron,  and  probably 
aid  in  its  carburisation.  The  molten  iron  is  drawn  off  at  intervals 
from  a  tap-hole  into  moulds,  and  is  known  as  cast  iron  or  pig  iron. 
The  slag  as  it  accumulates  overflows  in  a  regular  stream  through 
an  opening  known  as  the  slag  hole.  When  such  a  furnace  is  in  full 
blast,  fresh  charges  of  materials  are  introduced  at  regular  intervals, 
and  the  process  continues  uninterruptedly  for  years.  The  metal 
obtained  from  the  blast-furnace  is  far  from  pure  iron,  but  contains 
varying  quantities  of  carbon,  silicon,  phosphorus,  sulphur,  and 
manganese. 

The  carbon  may  be  present  either  in  combination  with  iron  as 
a  carbide,  or  distributed  throughout  the  metal  as  fine  particles  of 
graphite,  or  in  both  of  these  forms.  White  cast  iron  contains  its 
carbon  in  the  combined  form,  while  grey  cast  iroji  owes  its  grey 
colour  to  the  presence  of  minute  crystals  of  graphite  disseminated 
throughout  the  metal.  When  grey  cast  iron  is  dissolved  in  hydro- 
chloric acid,  the  graphite  remains  behind  as  a  black  powder  ;  but 
on  similarly  treating  iron  containing  combined  carbon,  the  carbon 
unites  with  the  hydrogen,  forming  various  hydrocarbons,  which 
impart  to  the  escaping  gas  a  characteristic  and  unpleasant  smell. 
Average  cast  iron  contains  from  90  to  95  per  cent,  of  iron,  and  3  to 
5  per  cent,  of  carbon.  Spiegel  is  a  variety  of  white  cast  iron  con- 
taining 3.5  to  6  per  cent,  of  carbon,  and  from  5  to  20  per  cent,  of 
manganese.  WTith  more  than  20  per  cent,  of  manganese,  the 
metal  is  termed  ferro-manganesc. 

Purification. — The  properties  of  iron  are  greatly  modified  by 
the  presence  of  various  impurities,  especially  carbon,  and  for 
different  purposes  for  which  iron  is  used,  metal  of  different  degrees 
of  purity  is  required.  The  purest  form  of  ordinary  commercial 


Iron  675 

iron  is  known  as  wrought  iron,  while  steel  is  intermediate  between 
this  and  cast  iron. 

The  process  by  which  cast  iron  is  converted  into  wrought  iron 
is  termed  puddling;  and  the  method  is  called  either  dry  puddling 
or pig-boiling,  depending  upon  whether  the  cast  iron  is  subjected 
to  a  preliminary  refining  or  not.  The  chemical  reactions  in  both 
cases  are  the  same,  and  consist  in  the  oxidation  of  the  impurities  ; 
the  carbon  being  expelled  as  carbon  dioxide,  while  the  oxides  of 
silicon,  phosphorus,  and  manganese  pass  into  the  slag.  The 
method  oi  pig-boiling  is  almost  exclusively  adopted. 

The  cast  iron  is  melted  in  a  reverberatory  furnace,  the  working 
bottom  of  which,  as  well  as  the  lining  (or  fettling),  consists  of  a 
layer  of  ferric  oxide.  The  decarburisation  of  the  iron  is  mainly 
effected  by  means  of  the  oxide  of  iron  derived  from  the  fettling  ; 
and  for  some  time  the  molten  mass  appears  to  boil,  owing  to  the 
escape  of  carbon  monoxide.  As  the  impurities  are  oxidised  and 
removed,  the  mass  becomes  pasty  (owing  to  the  fact  that  the 
melting-point  of  pure  iron  is  much  higher  than  that  of  cast  iron), 
and  is  then  worked  up  into  lumps,  or  blooms,  which  are  ultimately 
removed  and  placed  under  a  steam  hammer,  whereby  admixed  slag 
is  squeezed  out,  and  the  metal  is  welded  into  a  solid  mass. 

Wrought  iron  contains  from  0.06  to  0.15  per  cent,  of  carbon. 

Steel  may  be  produced  either  from  wrought  iron,  by  adding 
carbon,  or  from  cast  iron  by  removing  that  impurity.  Formerly 
steel  was  exclusively  obtained  by  the  first  method,  by  what  is 
known  as  the  cementation  process.  This  simply  consists  in  heating 
the  bars  of  iron,  buried  in  broken  charcoal,  for  several  days  to  a  red 
heat.  The  precise  nature  of  the  chemical  change  which  results  in 
the  carbunsation  of  the  iron  is  not  definitely  established.  In  all 
probability  the  carbon  is  conveyed  into  the  body  of  the  metal 
(which  is  not  even  heated  to  the  softening  point)  by  the  intervention 
of  iron  carbonyl  ;  the  carbon  monoxide  being  formed  by  the  union 
of  the  carbon  with  the  air  retained  within  the  layer  of  charcoal. 
At  the  conclusion  of  the  operation  the  iron  presents  a  blistered 
appearance,  and  on  this  account  is  termed  blister-steel. 

At  the  present  time  steel  is  mostly  produced  by  the  Bessemer 
process,  which  consists  in  oxidising  the  impurities  present  in  cast 
iron,  by  blowing  through  the  molten  metal  a  blast  of  air.  This 
operation  is  performed  in  a  large  pear-shaped  vessel,  known  as  a 
converter,  which  is  mounted  on  trunnions,  and  through  the  bottom 
of  which  a  powerful  air  blast  can  be  admitted.  The  converter  is 


6;6 


Inorganic  Chemistry 


tilted  into  a  horizontal  position,  and  a  quantity  of  molten  cast  iron 
is  run  in.  The  air  blast  is  then  started  and  the  converter  immedi- 
ately swung  back  into  a  vertical  position.  In  the  course  of  a  very 
short  time  the  whole,  of  the  impurities  are  burnt  away,  and  the 
stage  at  which  the  operation  is  complete  is  sharply  marked,  by  the 
sudden  disappearance  of  the  flame  from  the  open  mouth  of  the 
converter.  The  converter  is  once  more  swung  into  a  horizontal 
position,  and  the  blast  is  stopped.  The  exact  quantity  of  molten 
spiegel  is  then  added  to  supply  the  carbon  required  to  convert  the 
entire  charge  into  steel,  and  the  blast  is  turned  on  for  a  few 
moments  in  order  to  thoroughly  mix  the  materials,  after  which  the 
contents  are  poured  out  into  the  casting  ladle. 

The  comparative  purity  of  the  three  forms  of  iron  will  be  seen 
from  the  three  following  typical  examples  : — 

Cast  Iron. 


Carbon     . 

•     3-8i 

Silicon 

.     1.68 

Phosphorus 

.  '  0.70 

Sulphur    . 

.     0.60 

Manganese 

.     0.41 

7 

.20 

Iron 

92 

.80 

Steel. 

Wrought  Iron. 

0.65 

O.  IO 

0.07 

0.05 

0.03 

0.15 

0.02 

O.O5 

0.40 

0.07 



I.I7               •  0.42 

9^ 

3.83                           99.58 

IOO.OO 


IOO.OO 


Properties. — Pure  iron  is  a  white  lustrous  metal,  capable  of 
taking  a  high  polish.  Its  specific  gravity  is  7.84  to  8.139.  It  is 
more  difficultly  fusible  and  more  malleable  than  wrought  iron,  but 
at  a  red  heat  it  becomes  soft  and  can  be  welded.  The  physical 
properties  usually  associated  with  iron  are  in  reality  those  of 
iron  containing  varying  amounts  of  impurities  :  thus,  pure  iron 
when  rendered  magnetic  quickly  loses  this  property,  whereas 
steel  retains  its  magnetism  at  ordinary  temperatures,  losing  it, 
however,  when  heated.  Pure  iron,  when  heated  and  suddenly 
cooled,  does  not  take  a  temper,  while  steel  when  so  treated  be- 
comes extremely  hard  and  brittle. 

Iron  is  unacted  upon  by  dry  air  at  ordinary  temperatures, 
but  in  moist  air  it  quickly  becomes  coated  with  rust.  The  pre- 
sence of  the  atmospheric  carbon  dioxide  appears  to  be  indispensable 
to  the  rusting  of  iron,  the  first  stage  in  the  process  being  the  for- 


Oxides  of  Iron  677 

mation  of  ferrous  carbonate.*  Dilute  nitric  acid  dissolves  it, 
forming  ferrous  nitrate  and  ammonium  nitrate  ;  with  stronger 
nitric  acid,  ferric  nitrate  and  oxides  of  nitrogen  are  formed. 

Concentrated  nitric  acid  (specific  gravity,  1.45)  is  without  solvent 
action  upon  iron.  A  strip  of  iron  which  has  been  immersed  in 
such  strong  acid  is  unacted  upon  when  afterwards  dipped  into 
the  more  dilute  acid,  and  is  also  incapable  of  precipitating  metallic 
copper  from  a  solution  of  copper  sulphate.  Iron  in  this  condition 
is  said  to  ^passive.  Other  oxidising  agents,  as  chromic  acid,  or 
hydrogen  peroxide,  are  capable  of  bringing  about  the  same  result. 
It  is  believed  that  this  condition  is  due  to  the  formation  of  a  film 
of  the  oxide  Fe3O4  upon  the  surface. 

Finely  divided  iron  takes  fire  spontaneously  in  chlorine  ;  and 
when  gently  warmed  in  sulphur  dioxide  it  combines  with  that  gas 
with  incandescence.  It  absorbs  carbon  monoxide  with  formation 
of  iron  carbonyl,  Fe(CO)5.  When  heated  in  ammonia  it  forms  a 
nitride,  Fe4N2  (see  page  278). 

Oxides  of  Iron. — Three  oxides  of  iron  are  known,  namely  : — 

Ferrous  oxide  (iron  monoxide)      ,  FeO. 

Ferric  oxide  (iron  sesquioxide)      .         .     Fe2O3. 
Ferroso-ferric  oxide  (magnetic  oxide)    .     Fe3O4,  or  Fe2O3,FeO. 

The  two  first  are  basic  oxides,  giving  rise  respectively  to  ferrous 
and  ferric  salts  ;  the  third  yields  both  ferrous  and  ferric  salts. 

Ferric  oxide  combines  with  certain  more  basic  oxides,  form- 
ing compounds  analogous  to  Fe2O3,FeO  ;  such  as  Fe2O3,CaO, 
Fe2O3,ZnO.  These  are  known  asferrites. 

Ferrous  Oxide  (protoxide  of  iron),  FeO,  is  formed  as  an  inter- 
mediate product  during  the  reduction  of  ferric  oxide  by  hydrogen 
or  carbon  monoxide  ;  but  it  is  difficult  to  obtain  it  free  from  either 
the  higher  oxide  or  the  metal.  It  is  also  formed  when  ferrous 
oxalate  is  heated  out  of  contact  with  air.  It  is  a  black  powder, 
which  oxidises  in  the  air,  and  which  dissolves  in  acids  yielding 
ferrous  salts. 

Ferrous  Hydroxide,  Fe(HO)2,  or  FeO,H2O,  is  obtained  as  a 
white  precipitate  when  potassium  hydroxide  is  added  to  a  solution 
of  a  ferrous  salt  with  entire  exclusion  of  air.  In  the  presence  of 
air  it  is  green.  It  readily  absorbs  oxygen  and  passes  into  ferric 
oxide. 

*  Moody,  Jour.  Chem.  Soc.t  1906. 


678  Inorganic  Chemistry 

Ferric  Oxide  (sesquioxide  of  iron\  Fe2O3,  occurs  in  brilliant 
black  crystals  belonging  to  the  hexagonal  system,  in  specular 
iron  ore.  It  is  obtained  as  a  red  amorphous  powder  by  heating 
hydrated  ferric  oxide,  ferrous  sulphate,  or  ferrous  carbonate.  In 
a  crystalline  condition  it  may  be  produced  by  carefully  heating  a 
mixture  of  ferrous  sulphate  and  common  salt,  or  by  heating  the 
amorphous  oxide  in  gaseous  hydrochloric  acid.  The  natural  com- 
pound, and  also  the  artificial  substance  after  strong  ignition,  is 
only  slowly  dissolved  by  acids.  Ferric  oxide  is  extremely  hygro- 
scopic. When  strongly  heated  it  is  partially  converted  into  Fe3O4. 
The  amorphous  substance,  obtained  by  distilling  ferrous  sulphate 
for  the  manufacture  of  Nordhausen  sulphuric  acid,  is  employed  as 
a  red  pigment  and  a  polishing  powder  under  the  name  of  rouge. 

Ferric  Hydroxide,  or  Hydrated  Ferric  Oxide,  Fe2(HO)6,  or 
Fe2O3,3H2O. — When  an  excess  of  ammonia  is  added  to  a  solution 
of  ferric  chloride,  and  the  voluminous  brown  precipitate  is  dried 
at  a  moderate  temperature,  it  has  the  composition  Fe2O3,3H2O. 
On  exposure  to  various  temperatures,  or  by  precipitation  under 
various  conditions,  hydrated  oxides  of  the  composition  Fe2O3, 
2H2O  ;  Fe2O3,H2O,  and  others,  have  been  obtained  ;  and  several 
of  these  compounds  occur  in  nature.  Ordinary  rust  of  iron  consists 
of  a  mixture  of  hydrated  ferric  and  ferrous  oxides  with  ferrous 
carbonate,  in  varying  proportions  depending  upon  conditions. 
Exposure  to  air  gradually  oxidises  the  ferrous  oxide  and  carbonate 
into  ferric  oxide. 

The  monohydrate  Fe2O3,H2O  has  been  obtained  as  a  soluble  modification, 
by  heating  an  acetic  acid  solution  of  precipitated  ferric  hydroxide  to  100°  in 
sealed  vessels.  On  the  addition  of  sulphuric  acid,  a  brown  precipitate  is 
obtained,  having  the  composition  Fe2O3,H2O,  which  is  insoluble  in  acids, 
but  soluble  in  water.  The  solution  gives  no  reaction  with  potassium  ferro- 
cyanide.  Another  soluble  hydroxide  is  produced  by  dissolving  the  ordinary 
precipitated  hydroxide  in  ferric  chloride,  and  subjecting  the  solution  to 
dialysis.  This  solution  is  employed  in  medicine  under  the  name  of  dialysed 
it  on. 

Ferroso-ferrie  Oxide,  Fe3O4,  occurs  native  as  magnetite  and 
magnetic  oxide  of  iron',  the  magnetic  variety  being  known  also 
as  loadstone.  When  iron  is  heated  in  the  air,  the  black  film 
which  forms  (the  so-called  iron-scale  or  hammer-scale}  consists  of 
the  oxide  Fe3O4,  with  more  or  less  ferric  oxide,  Fe2O3,  upon  the 
outer  surface.  It  is  also  produced  when  steam  or  carbon  dioxide 
is  passed  over  heated  iron,  with  evolution  of  hydrogen  and  carbon 
monoxide  respectively,  these  reactions  being  the  reverse  of  those  by 
which  oxides  of  iron  are  reduced  by  hydrogen  or  carbon  monoxide. 


Ferrous  Sulphate  679 

This  oxide  is  also  formed  as  a  black  precipitate  when  ammonia 
is  added  to  a  solution  containing  mixed  ferrous  and  ferric  salts,  and 
the  mixture  gently  warmed. 

Ferrates.  —  These  compounds  correspond  to  the  manganates, 
but  neither  the  acid  H2FeO4  nor  the  oxide  FeO3  are  known. 
Potassium  ferrate,  K2FeO4,  is  formed  when  chlorine  is  passed 
through  a  solution  of  potassium  hydroxide  in  which  ferric  hydroxide 
is  suspended. 

FERROUS   SALTS. 

Ferrous  Chloride,  FeCl2.—  The  anhydrous  compound  is  ob- 
tained by  heating  iron  wire  in  gaseous  hydrochloric  acid,  when  the 
salt  sublimes  in  the  form  of  white  deliquescent  crystals.  In  aqueous 
solution  it  is  obtained  when  iron  is  dissolved  in  hydrochloric  acid, 
and  is  deposited  in  pale  blue-green  crystals  of  FeCl2,4H2O. 

When  heated  in  the  air  it  is  converted  into  ferric  oxide  and 
chloride,  the  latter  volatilising  — 


When  volatilised  in  an  atmosphere  of  hydrochloric  acid  its 
vapour-density  at  high  temperatures  corresponds  to  the  formula 
FeCl2  ;  at  lower  temperatures  it  lies  between  the  values  required 
for  FeCl2  and  Fe2Cl4. 

When  strongly  heated  in  a  current  of  steam  it  is  decomposed  as 
follows  — 


Ferrous  Sulphate  (green  vitriot),  FeSO4,7H2O,  is  obtained 
when  iron  is  dissolved  in  sulphuric  acid.  It  is  prepared  on  a 
large  scale  by  exposing  heaps  of  iron  pyrites,  FeS2,  to  the  action 
of  air  and  moisture.  The  liquor  which  drains  away  contains 
ferrous  sulphate  and  sulphuric  acid,  and  the  latter  is  converted 
into  ferrous  sulphate  by  the  introduction  of  scrap  iron. 

Ferrous  sulphate  forms  pale  green  monosymmetric  crystals, 
which  effloresce  on  exposure  to  the  air.  They  are  soluble  in  water 
to  the  extent  of  70  parts  in  100  parts  of  water  at  15°,  and  370  parts 
in  loo  parts  at  90°.  At  100°  the  crystals  lose  6H2O,  being  con- 
verted into  FeSO4,H2O. 

If  a  crystal  of  zinc  sulphate  be  thrown  into  a  supersaturated 
solution  of  ferrous  sulphate,  the  iron  salt  is  deposited  in  rhombic 


680  Inorganic  Chemistry 

prisms  (isomorphous  with  zinc  sulphate).  On  the  other  hand,  if  a 
crystal  of  copper  sulphate  be  added,  asymmetric  (triclinic)  crystals 
of  FeSO4,5H2O  (isomorphous  with  copper  sulphate)  are  formed. 

Ferrous  sulphate  forms  double  salts  with  the  sulphates  of  the 
alkalies.  Thus,  when  mixed  with  ammonium  sulphate  in  the  re- 
quisite proportions,  ammonium  ferrous  sulphate,  FeSO4,(NH4)2SO4, 
6H2O,  is  obtained.  This  salt  is  less  readily  oxidised  on  exposure 
to  air  than  ferrous  sulphate  itself. 

Ferrous  salts  give,  with  potassium  ferrocyanide  (K4Fe(CN)c, 
or  4KCN,Fe(CN)2),  a  white  precipitate  of  potassium  ferrous  ferro- 
cyanide (FeK2Fe(CN)6,  or  2KCN,2Fe(CN)2).  The  precipitate  is 
quickly  oxidised,  and  becomes  blue.  With  potassium  ferricyanide 
(K3Fe(CN)6,  or  3KCN;Fe(CN)3),  ferrous  salts  yield  a  blue  pre- 
cipitate of  ferrous  ferricyanide  (T^trnbul^s  blue}  (Fe3{Fe(CN)6!o,  or 
3Fe(CN)2,2Fe(CN)3),  thus— 

=  Fe3{Fe(CN)6}2 


FERRIC   SALTS. 

Ferric  Chloride,  FeCl3,  is  prepared  in  the  anhydrous  state  by 
passing  dry  chlorine  over  heated  iron  wire.  In  solution  it  may 
be  obtained  by  dissolving  iron  in  aqua  regia;  or  ferric  oxide  in 
hydrochloric  acid.  The  anhydrous  salt  forms  nearly  black  crystals, 
appearing  deep  red  by  transmitted  light.  It  readily  volatilises,  and 
at  temperatures  above  700°  the  density  of  its  vapour  corresponds  to 
the  formula  FeCl3,  while  at  lower  temperatures  its  density  agrees 
more  nearly  with  the  formula  Fe2Cl6. 

Ferric  chloride  is  extremely  deliquescent,  and  readily  dissolves  in 
water.  When  the  solution  is  slowly  evaporated,  yellow  crystals  are 
deposited,  having  the  composition  Fe2Cl6,12H2O  (or  FeCl3,6H2O). 
When  a  dilute  solution  of  ferric  chloride  is  boiled,  it  decomposes, 
forming  either  an  insoluble  oxychloride  or  a  soluble  hydroxide  and 
free  hydrochloric  acid  (depending  upon  the  strength  of  the  solution). 

Ferric  Sulphate,  Fe2(SO4)3,  is  prepared  by  the  addition  of  sul- 
phuric or  nitric  acids  to  a  solution  of  ferrous  sulphate  — 


The  brown  solution,  on  evaporation,  leaves  the  anhydrous  salt 
as  a  white  mass.     When  the  requisite  quantity  of  potassium  sul- 


Sulphides  of  Iron  68  1 

phate  is  dissolved  in  a  strong  solution  of  ferric  sulphate  at  o°, 
the  double  potassium  iron  sulphate  (iron  alum),  K2SO4,Fe2(SO4)3, 
24H2O,  separates  out  in  the  form  of  violet  octahedrons. 

Ferric  salts  give,  with  potassium  ferrocyanide  (K4Fe(CN)6,  or 
4KCN,Fe(CN)2),  a  dark  blue  precipitate  of  ferric  ferrocyanide 
(Prussian  blue],  4Fe(CN)3,3Fe(CN),)  or  Fe4{Fe(CN)0}3— 

4FeCl3  +  3K4Fe(C'N)6=Fe4{Fe(CN)6}3  +  12KCL 
With  potassium  ferricyanide  ferric  salts  give  no  precipitate. 

SULPHIDES  OF  IRON. 

Ferrous  Sulphide,  FeS.  —  When  a  white-hot  bar  of  wrought 
iron  is  dipped  into  melted  sulphur,  the  elements  unite  ;  and  the 
readily  fusible  monosulphide  of  iron  falls  to  the  bottom.  It  may 
be  prepared  by  throwing  into  a  red-hot  crucible  a  mixture  of  iron 
filings  and  sulphur.  So  obtained,  it  is  a  dark,  yellowish-grey, 
metallic-looking  mass.  When  heated  out  of  contact  with  air,  it 
does  not  part  with  sulphur,  but  in  the  presence  of  air  is  converted 
into  ferric  oxide  and  sulphur  dioxide.  Ferrous  sulphide  is  pre- 
cipitated from  either  ferrous  or  ferric  solutions,  by  alkaline  sul- 
phides, as  a  black  amorphous  powder,  which  in  the  moist  state  is 
quickly  oxidised  by  the  air.  Dilute  sulphuric  acid,  or  hydrochloric 
acid,  decomposes  ferrous  sulphide,  with  evolution  of  sulphuretted 
hydrogen. 

Iron  Sesquisulphide,  Fe2S3,  is  formed  when  equal  weights  of 
iron  and  sulphur  are  heated  to  a  moderate  temperature.  It  can- 
not be  obtained  by  precipitation  from  a  ferric  salt,  as  the  product 
so  formed  consists  of  ferrous  sulphide  and  sulphur  — 


It  is  a  yellow,  metallic-looking  solid,  which  is  decomposed  by 
dilute  hydrochloric  acid,  yielding  sulphuretted  hydrogen. 

Ferric  Bisulphide,  FeS2,  occurs  in  nature  in  large  quantities  as 
iron  pyrites,  sometimes  in  the  massive  condition,  and  at  others  in 
the  form  of  brass-yellow  cubical  crystals.  In  many  cases  the 
native  compound  bears  the  impression,  or  assumes  the  shape, 
of  various  organised  forms,  such  as  wood,  ammonites,  £c.,  the 
mineral  having  been  formed  by  the  reducing  action  of  the  organic 
matter  upon  ferrous  sulphate  in  solution.  Ferric  disulphide  is 
also  found  in  the  form  of  brass-like,  rhombic  crystals  in  radiated 
pyrites. 


682  Inorganic  Chemistry 

The  compound  may  be  prepared  by  heating  to  a  low  red  heat 
a  mixture  of  ferrous  sulphide  and  sulphur. 

Ferric  disulphide  is  unacted  upon  by  dilute  acids  :  hot  con- 
centrated hydrochloric  acid  decomposes  it,  with  liberation  of  sul- 
phur and  sulphuretted  hydrogen.  When  heated  in  hydrogen, 
sulphur  is  evolved  (which  partly  combines  with  the  hydrogen), 
and  ferrous  sulphide  remains.  When  heated  in  the  air,  ferric 
oxide  and  sulphur  dioxide  are  formed. 

Ferroso-ferrie  Sulphide  (magnetic  pyrites),  Fe3S45  occurs  in 
the  form  of  hexagonal  crystals.  Like  the  corresponding  oxide, 
this  compound  is  attracted  by  the  magnet,  and  is  itself  sometimes 
magnetic. 

COBALT. 
Symbol,  Co.     Atomic  vreight= 58. 97. 

Occurrence. — Cobalt  is  not  found  uncombined  in  nature.  Its 
chief  natural  compounds,  which  are  only  sparsely  distributed,  are 
speis's-cobalt,  or  smaltine,  CoAs2  ;  cobalt  glance,  CoAsS,  in  both  of 
which  the  cobalt  is  partially  replaced  by  nickel  and  iron  ;  and 
cobalt-bloom,  Co3(AsO4)2,8H2O. 

Modes  Of  Formation.— Cobalt  is  obtained  by  reducing  the 
oxide,  or  the  chloride,  in  a  stream  of  hydrogen,  or  by  strongly 
heating  cobalt  oxalate  in  a  closed  crucible.  .It  is  also  readily 
obtained  by  reduction  of  its  oxide  with  powdered  aluminium. 

Properties. — Cobalt  is  an  almost  white,  hard  metal,  which, 
when  polished,  resembles  nickel,  but  is  slightly  bluer.  It  is 
malleable,  and  when  heated  is  very  ductile.  Like  both  iron  and 
nickel,  it  is  attracted  by  the  magnet ;  but  unlike  these,  it  retains 
this  property,  even  at  a  red  heat.  In  the  massive  form,  cobalt  is 
unacted  upon  by  the  air  ;  but  the  finely-powdered  metal,  obtained 
by  the  reduction  of  the  oxide  in  hydrogen,  rapidly  oxidises  on 
exposure  to  the  air,  sometimes  with  incandescence.  When  heated 
in  the  air,  it  forms  the  oxide  Co3O4.  Cobalt  decomposes  steam  at 
a  red  heat,  yielding  cobaltous  oxide,  CoO. 

Oxides  Of  Cobalt. — Three  oxides  of  cobalt  are  recognised, 
namely,  cobaltous  oxide,  CoO  ;  cobaltic  oxide,  Co2O3 ;  and  cobalto- 
cobaltic  oxide,  Co3O4. 

Four  other  oxides  are  known,  which  are  regarded  as  compounds  of  the  two 
first,  having  the  composition  2CoO,Co2O3;  3CoO,Co2O3 ;  4CoO,Co2O3.' 
(iCoO,Co2O3. 


Cobaltous  Chloride  683 

The  monoxide,  CoO,  is  basic,  and  yields  the  cobaltous  salts. 
The  sesquioxide,  Co2O3,  is  feebly  basic,  forming  only  unstable 
salts.  Stable  double  salts,  however,  corresponding  to  this  oxide 
are  known. 

Cobaltous  Oxide  (cobalt  monoxide],  CoO,  is  formed  when  the 
sesquioxide  is  heated  to  redness  in  a  stream  of  carbon  dioxide,  or 
gently  heated  in  hydrogen.  It  is  also  obtained  when  the  carbo- 
nate or  hydroxide  is  heated  in  the  absence  of  air.  It  forms  a  drab- 
coloured  powder,  which  is  unacted  upon  by  the  air,  but  when  heated, 
forms  Co3O4.  When  heated  in  either  hydrogen  or  carbon  mon- 
oxide, it  is  reduced  to  metallic  cobalt. 

Cobaltous  Hydroxide,  Co(HO)2.— When  potassium  hydroxide 
is  added  to  a  solution  of  a  cobaltous  salt,  a  blue  basic  hydrate  is 
precipitated,  which,  on  boiling,  is  converted  into  the  pink  hydroxide 
Co(HO)2.  It  turns  brown  on  exposure  to  the  air,  by  the  absorp- 
tion of  oxygen.  Both  the  oxide  and  hydroxide  are  really  soluble 
in  acids,  giving  cobaltous  salts. 

Cobaltie  Oxide  (cobalt  sesquioxide},  Co2O3,  is  obtained  by  care- 
fully heating  cobaltous  nitrate  until  red  fumes  cease  to  be  evolved. 
It  is  a  dark  grey  powder,  which,  when  strongly  heated,  is  con- 
verted into  the  intermediate  black  oxide,  Co3O4.  Cobaltie  oxide 
dissolves  in  cold  acids,  forming  brown  solutions,  which  contain 
unstable  cobaltic  salts.  When  warmed,  these  are  converted  into 
cobaltous  salts,  with  evolution  of  oxygen  in  the  case  of  oxy-salts, 
and  of  the  halogen  from  haloid  salts.  This  sesquioxide,  therefore, 
behaves  as  a  peroxide. 

Cobaltie  Hydroxide,  Co2(HO)6,  orCo2O3,3H2O,is  obtained  as  a 
nearly  black  precipitate,  by  the  addition  of  an  alkaline  hypochlorite 
to  a  cobaltous  salt.  With  acids  it  behaves  as  the  oxide. 

Cobalto-Cobaltie  Oxide,  Co3O4,  is  formed  as  a  black  powder, 
when  the  sesquioxide  is  strongly  heated  in  air. 


COBALTOUS  SALTS. 

Cobaltous  Chloride,  CoCl2. — When  the  carbonate,  or  any  of 
the  oxides,  are  dissolved  in  hydrochloric  acid,  the  concentrated 
solution  deposits  dark  red  prisms  (monosymmetric),  having  the 
composition  CoCl2,6H2O.  When  exposed  over  sulphuric  acid,  they 
lose  4H2O,  and  are  converted  into  a  rose-red  salt,  CoCl2,2H2O, 
which  reabsorbs  moisture  from  the  air  to  form  the  hexahydrate. 


684  Inorganic  Chemistry 

When  the  dihydrate  is  heated  to  about  100°,  it  is  converted  into 
violet-blue  crystals  of  CoCl2,H2O  ;  and  at  120°  it  becomes  an- 
hydrous, and  is  blue.  The  blue  salts,  on  exposure  to  the  air, 
rapidly  rehydrate  themselves,  and  become  pink. 

Cobaltous  chloride  dissolves  in  alcohol,  giving  a  deep  blue  solu- 
tion, which,  on  the  addition  of  water,  also  becomes  pink.  This 
property  of  forming  pink  hydrated  salts,  which  become  blue  or 
green  when  nearly  or  quite  anhydrous,  is  common  to  most  cobal- 
tous  salts.  Thus,  the  iodide  CoI2,6H2O  forms  rose-coloured 
crystals.  When  gently  heated,  it  changes  to  a  moss-green  salt, 
CoI2,2H2O,  which,  when  dehydrated,  becomes  nearly  black. 

CobaltOUS  Sulphate,  CoSO4,7H2O,  is  obtained  by  dissolving 
the  carbonate  or  oxides  in  sulphuric  acid,  and  is  deposited  from 
the  solution  in  dark  red  crystals,  isomorphous  with  ferrous  sul- 
phate. Cobalt  sulphate,  like  the  sulphates  of  iron  and  nickel., 
forms  double  salts  with  alkaline  sulphates,  of  which  cobalt  potas- 
sium sulphate,  CoSO4,K2SO4,6H2O,  is  the  best  known. 

Cobaltie  Salts. — Single  salts  corresponding  to  cobalt  sesqui- 
oxide  are  unstable,  and  exist  only  in  solution.  More  stable  double 
salts  are  known.  Thus,  when  potassium  nitrite  is  added  to  an 
acetic  acid  solution  of  cobalt  chloride,  a  yellow  crystalline  precipi- 
tate' is  obtained,  consisting  of  the  double  nitrite  of  cobalt  and 
potassium — 

2CoCl2  +  10KNO2  +  4HNO2-Co2(NO2)6,6KNO2  + 
2NO  +  4KC!4-2H2O. 

The  formation  of  this  compound  is  made  use  of  for  separating  cobalt  from 
nickel,  the  latter  element  yielding  no  corresponding  double  nitrite.  In  the 
presence,  however,  of  salts  of  barium,  strontium,  or  calcium,  nickel  forms, 
with  potassium  nitrite,  triple  salts,  such  as  Ni(NO.2)2>Ba(NO.2)2l2KNO2,  which 
are  precipitated  as  yellow  crystalline  powders.  Hence,  in  the  presence  of 
metals  of  the  alkaline  earths,  nickel  and  cobalt  cannot  be  separated  by  this 
method. 

SULPHIDES  OF  COBALT. 

CobaltOUS  Sulphide,  CoS,  is  obtained  by  heating  cobaltous 
oxide  with  sulphur,  or  by  fusing  a  mixture  of  cobalt  sulphate, 
barium  sulphide,  and  common  salt.  It  forms  bronze-coloured 
crystals,  which  are  soluble  in  strong  hydrochloric  acid.  Cobalt 
sulphide  is  precipitated  as  a  black  amorphous  powder  when 
ammonium  sulphide  is  added  to  a  cobalt  solution.  The  precipi- 
tate slowly  dissolves  in  dilute  mineral  acids,  but  is  insoluble  in 


Cobaltamines  685 

acetic  acid.  When  heated  in  a  stream  of  sulphuretted  hydrogen, 
it  yields  the  sesquisulphide  Co2S3  ;  and  if  mixed  with  sulphur,  and 
heated  in  a  current  of  hydrogen,  it  forms  the  disulphide  CoS2. 

Cobaltamines  (am maniacal  cobalt  compounds*}.  Cobalt  forms 
a  large  number  of  complex  ammoniacal  salts.  A  few  of  these 
contain  the  metal  in  the  divalent  condition,  and  are  known  as 
ammonio-cobaltous  salts;  but  by  far  the  larger  number  contain 
the  hexavalent  double  atom  Co2,  and  are  termed  ammonio-cobaltic 
compounds.  These  compounds  are  classified  as  follows  t  : — 

Ammonio-Cobaltous  Salts  are  formed  by  the  absorption  of  gaseous  am- 
mor.ia  by  anhydrous  cobaltous  salts,  or  by  dissolving  the  salts  in  strong 
aqueous  ammonia,  with  exclusion  of  air.  In  this  way  the  following  salts  have 
been  obtained — 

^  r-i    /.XTTT      f  which,  at  120°,  is  converted 
Ammomo-cobaltous  chloride,  CoCl2,6NH3    |     into  CoCL^NH,. 

Ammonio-cobaltous  sulphate,  CoSO4,6NH3. 
Ammonio-cobaltous  nitrate,  Co(NO3)2,6NH3,2H2O. 

Ammonio-Cobaltic  Salts.— These  may  be  arranged  under  the  following 
classes  and  subdivisions  : — 

I.  Hexammonio  Salts. — General  formula,  Co2(NH3)6'R6,  where  R  equals 
a  monacid  radical,  or  its  equivalent  of  di  or  tri  acid  radicals. 

f  Hexammonio-cobaltic  chloride  (dichro-cobaltic  chloride) 
Examples  \       Co2-(NH3)6'Cl6,2H2O. 

I.  Hexammonio-cobaltic  sulphate,  Co2<(NH3)6<(SO4)3,6H2O. 

II.   Octammonio  Salts — 

(a.)  Praseo%  Salts.—  General  formula,  Co2'(NH3)8'R6. 

-cobaltic  chloride,  Co2(NH3)8-Cl6,2H2O. 
Examples  \  Praseo-cobaltic  chloro-nitrate,  Co2(NH3)8-Cl4-(NOs)2, 


f  Praseo-cobaltic 
-j  Praseo-cobaltic 
I  2H20. 


(j8.)  Fusco  Salts.— General  formula,  Co2(NH3)s(HO)2-R4. 

C  Fusco-cobaltic  chloride,  Co2(NH3)8(HO)2-Cl4,2H2O. 
Examples  4  Fusco  -  cobaltic    sulphate,   Co2(NH3)s(HO)2'(SO4)2, 
I      2H2O. 

*  For  details  respecting  the  preparation  and  properties  of  these  salts  the 
student  is  referred  to  larger  works. 

f  On  the  constitution  of  metallammonium  compounds  generally,  see  Werner, 
Zcitschrift  fur  Anorganische  Chemie,  1893,  vol.  iii. 

t  These  names  denote  the  characteristic  colours  of  the  salts,  thus— prasi rnus, 
leek-green  ;  fuscus,  swarthy  ;  crocus,  yellow,  &c. 


686  Inorganic  Chemistry 

(7.)  Croceo  Salts.—  General  formula,  Co2(NH3)8(NO2)4'R2. 

Evamtles   /  Croceo-cobaltic  chloride,  Co2(NH8)8(NOa)4-Cl2. 
I.  Croceo-cobaltic  sulphate,  Co2(NH3)8(NO2)4'SO4. 

III.  Decammonio  Salts — 

(a.)  Roseo  Salts.—  General  formula,  Co2(NH3)10(H2O)2Re. 

Roseo-cobaltic  chloride,  Co2(NH3)10(H2O)2Cl6. 


Examples  -j  Roseo-cobaltic    sulphate,    Co2(NH3)10(H2O)2-(SO4)3t 
3H2O. 

Purpureo  Salts.—  General  formula,  Co2(NH3)10X2R4 

(where  X  and  R  are  either  the  same  or  different  acid  radicals). 

'Chloro-purpureo-cobalticchloride,Co2(NH3)i0Cl2'Cl4. 
Chloro-purpureo-cobaltic  sulphate,  Co2(NH3)10Cl2* 
Examples  \      (SO4)2. 

Bromo  -  purpureo  -  cobaltic    nitrate,    Co.,(NH3)10Br2' 
(NO,)* 

(-y.)  Xantho  Salts. — General  formula,  Co2(NH3)10(NO2)2'R4. 

(  Xantho-cobaltic  chloride,  Co2(NH3)10(NO2)2'Cl4. 
Examples  -!  Xantho-cobaltic    bromo-nitrate,    Co2(NH3)10(NOo)2- 
I      Br2'(N03)2. 

IV.  Oxy-decammonio  Salts.—  General  formula,  Co2(NH3)10R4-X'O(HO) 

(where  X  is  either  (HO)  or  an  acid  radical  either  the  same  as, 
or  different  from,  R). 

/'Oxy-decammonio  cobaltic  chloride,  Co0(NH3)10Cl4* 

Examples  \      (HO)-O'(HO). 

j  Anhydro  -  oxy  -  decammomo    cobalt    chloride, 
I     Co,(NH3)10Cl4-Cl2-0-(HO). 

V.  Dodecammonio  Salts  (luteo-cobaltic  salts).  — General  formula, Co2(HN3)12R6. 

E,         ,r     /  Luteo-cobaltic  chloride,  Co2(NH3)12Cl6. 

\Luteo-cobalticsulphate,  Co2(NH3)12(SO4)3,5H2O. 

When  cobalt  compounds  are  fused  with  borax,  a  clear  blue 
vitreous  mass  is  obtained,  which  contains  a  borate  of  cobalt.  A 
similar  blue  colour  is  imparted  to  ordinary  potash  glass  when  a 
small  quantity  of  a  cobalt  salt  is  added  to  the  molten  material, 
owing  to  the  formation  of  a  silicate  of  cobalt.  Under  the  name 
of  smalt,  this  substance  has-been  manufactured  for  use  as  a  pig- 
ment, by  fusing  the  roasted  cobalt  ore  with  quartz  sand  and  pearl- 
ash.  The  fused  mass  of  deep  blue  glass  is  then  finely  ground 
beneath  water. 


Nickel  Alloys  687 

NICKEL. 

Symbol,  Ni.     Atomic  weight  =  58. 7. 

Occurrence. — Nickel  occurs  chiefly  in  combination  with  arsenic 
as  kupfer  nickel*  Ni2As2 ;  white  nickel,  NiAs2 ;  nickel  glance, 
Ni2(AsS)2,  also  as  nickel  blende,  NiS.  Nickel  ore  almost  invari- 
ably contains  cobalt,  and  frequently  antimony  and  bismuth. 

Modes  Of  Formation.— Nickel  is  obtained  by  reducing  the 
oxide  with  carbon  at  a  high  temperature.  It  may  be  obtained  as  a 
black  powder  by  reducing  nickelous  oxide  in  a  stream  of  hydrogen, 
or  by  heating  nickelous  oxalate  out  of  contact  with  air.  It  is  also 
obtained  as  a  lustrous  coherent  deposit  by  the  electrolysis  of  an 
ammoniacal  solution  of  the  double  sulphate  of  nickel  and  ammonia. 

Nickel  in  a  high  state  of  purity  is  now  being  made  on  a  com- 
mercial scale  by  what  is  known  as  the  "  Mond's  "  process.  This 
consists  in  passing  carbon  monoxide  over  gently  heated  nickel 
oxide,  whereby  the  nickel  is  first  reduced  and  is  then  taken  up  by 
the  carbon  monoxide  to  form  nickel  carbonyl,  Ni(CO)4  (see  p.  299). 
This  volatile  compound  is  then  passed  through  tubes  which  are 
more  strongly  heated,  which  causes'  the  compound  to  decompose 
into  carbon  monoxide  (which  can  be  again  utilised)  and  metallic 
nickel.  In  this  way  the  metal  is  deposited  in  the  form  of  a  coherent 
solid,  entirely  free  from  cobalt,  with  which  nickel  is  always  associated 
in  its  ores. 

Properties. — Nickel  is  a  lustrous  white  metal,  with  a  faint 
yellow  tinge  when  compared  with  silver.  It  is  ductile  and  malle- 
able, and  at  the  same  time  very  hard  and  tenacious.  It  is  sus- 
ceptible of  a  very  high  polish.  Nickel  is  attracted  by  the  magnet, 
but  loses  this  property -when  moderately  heated.  When  obtained 
by  reduction  with  charcoal,  the  metal  contains  a  certain  amount  of 
carbon  (like  cast  iron),  which  renders  it  less  malleable,  and  when 
produced  by  reduction  of  the  oxalate  at  a  low  temperature  the 
powder  is  pyrophoric. 

In  the  massive  form,  nickel  is  unacted  upon  by  moderately  dry 
air,  but  in  moist  air  it  tarnishes,  and  becomes  covered  with  a  film 

*  Kupfer  nickel  signifies  the  false  copper,  and  was  applied  by  the  Germans 
in  the  Middle  Ages  to  this  ore,  which  resembled  a  copper  ore,  because  they 
tried  in  vain  to  extract  copper  from  it.  It  is  probable  that  this  ore  had  been 
smelted  along  with  copper  ores,  under  the  belief  that  it  contained  copper,  by 
the  early  ancients.  Thus,  a  coin,  235  B.C.,  has  been  found  to  contain  20  per 
cent,  of  nickel. 


688  Inorganic  Chemistry 

of  nickelous  oxide.  It  decomposes  steam  only  slowly  at  a  red 
heat,  and  is  slowly  attacked  by  dilute  hydrochloric  or  sulphuric 
acid  (contrast  iron). 

Nickel  is  largely  used  for  electro-plating  iron  and  steel  articles. 
It  is  also  employed  on  an  extensive  scale  in  the  production  of  nickel 
steel  for  modern  armour  plate,  in  which  the  proportion  of  nickel 
reaches  20  or  even  30  per  cent. 

Nickel  Alloys.  —  With  copper,  and  with  copper  and  zinc,  nickel 
furnishes  several  important  alloys.  The  small  coinage  in  use  in 
Belgium,  Germany,  and  the  United  States  consists  of  i  part  of 
nickel  and  3  parts  of  copper  ;  while  the  so-called  German  silver, 
or  nickel-silver,  contains  in  addition  about  1.5  part  of  zinc. 

Oxides  of  Nickel,  —  Three  oxides  of  nickel  have  been  obtained, 
namely,  nickelous  oxide,  NiO  ;  nickelic  oxide,  Ni2O3  ;  and  nickelo- 
nickelic  oxide,  Ni3O4.  The  first  alone  is  basic. 

Niekelous  Oxide  (nickel  monoxide},  NiO,  is  obtained  as  a 
greenish  powder  by  heating  nickel  carbonate  or  hydroxide  out  of 
contact  with  air.  It  is  dissolved  by  acids  yielding  nickel  salts. 
When  heated  in  hydrogen  or  carbon  monoxide  it  is  readily  re- 
duced to  the  metallic  state. 

Niekelous  Hydroxide,  Ni(HO)2,  is  obtained  in  a  pale  green 
precipitate  when  potassium  hydroxide  is  added  to  a  solution  of  a 
nickel  salt  ;  the  precipitate  has  the  composition  4Ni(HO)2,H2O. 
When  strongly  heated  it  is  converted  into  nickelous  oxide  and 
water.  It  is  readily  soluble  in  acids,  forming  the  nickel  salts,  and 
it  also  dissolves  in  ammonia  and  in  solutions  of  ammonium  salts. 

Nickel  Sesquioxide,  Ni2O3,  is  obtained  as  a  black  powder 
when  the  nitrate  is  decomposed  by  heat  at  the  lowest  temperature. 
With  hydrochloric  acid  and  sulphuric  acid  it  behaves  like  a  per- 
oxide, yielding  nickel  salts,  with  the  elimination  of  chlorine  and 
oxygen  respectively  — 

Ni203  +  6HQ     =  2NiCl2  +3H20  +  C12. 


It  is  soluble  in  ammonia,  with  evolution  of  nitrogen  — 


Hydrated  Sesquioxide  of  Nickel,  Ni2(HO6),  or  Ni2O3,3H2O. 
When  chlorine  is  passed  through  water  or  sodium  hydroxide,  in  which 
nickelous  hydroxide,  Ni(HO)2,  is  suspended,  a  black  powder  is  ob- 
tained having  the  composition  Ni2O3,3H2O.  The  same  compound 
is  obtained  when  a  nickel  salt  is  added  to  a  solution  of  bleaching- 
powder.  In  contact  with  acids  and  ammonia  it  behaveslike  the  oxide. 


Nickelous  Sulphide  689 

Nickelo-niekelic  Oxide,  Ni3O4,  is  obtained  as  a  grey  metallic- 
looking  mass,  when  nickel  chloride  is  heated  to  about  400°  in  a 
stream  of  oxygen. 

Nickel  Salts. — Nickel  forms  only  one  series  of  salts,  corre- 
sponding to  the  monoxide.  In  the  anhydrous  state  these  are 
usually  yellowish,  while  in  the  hydrated  condition  they  are  green. 

Nickel  Chloride,  NiCU,  is  obtained  as  a  yellow  amorphous 
mass,  by  dissolving  the  oxide  or  carbonate  in  hydrochloric  acid, 
and  evaporating  the  solution  to  dryness.  When  heated  in  a 
current  of  chlorine  it  sublimes  in  the  form  of  lustrous  golden 
yellow  scales,  which  dissolve  in  water,  forming  a  green  solution. 
From  the  aqueous  solution,  green  crystals  of  the  composition 
NiCl2,6H2O  are  deposited. 

Anhydrous  nickel  chloride  absorbs  gaseous  ammonia,  forming  the 
compound  NiCl2,6NH3,  which  when  deposited  from  an  aqueous 
solution  forms  blue  octahedrons. 

Nickel  Sulphate,  NiSO4,7H2O,  is  produced  when  the  metal, 
the  carbonate,  or  the  oxide  is  dissolved  in  dilute  sulphuric  acid, 
and  the  concentrated  solution  is  allowed  to  crystallise  at  the  ordi- 
nary temperature.  It  forms  green  crystals,  isomorphous  with 
magnesium  sulphate.  When  heated  to  100°  the  crystals  lose 
6H2O,  and  above  300°  the  salt  becomes  anhydrous.  The  anhy- 
drous salt  absorbs  gaseous  ammonia,  being  converted  into  a  pale 
violet  powder  having  the  composition  NiSO4,6NH3.  When  nickel 
sulphate  is  dissolved  in  strong  aqueous  ammonia,  the  solution 
deposits  dark  blue  tetragonal  crystals  of  NiSO,4,4NH3,2H2O. 

With  sulphates  of  the  alkalies,  nickel  sulphate  forms  double 
salts,  of  which  the  ammonium  salt  is  the  most  important,  NiSO4, 
(NH4)2SO4,6H2O.  It  is  obtained  by  mixing  concentrated  solu- 
tions of  the  two  sulphates  in  the  requisite  proportions.  This  salt 
is  employed  in  the  process  of  nickel-plating. 

Niekelous  Sulphide  (nickel  monosulphide],  NiS,  occurs  as  the 
mineral  capillary  pyrites.  It  is  obtained  as  a  bronze-like  mass, 
insoluble  in  hydrochloric  acid,  by  heating  sulphur  and  nickel 
together.  In  the  hydrated  condition  nickel  sulphide  is  precipitated 
as  an  amorphous  black  powder,  on  the  addition  of  ammonium 
sulphide  to  a  nickel  salt.  The  precipitate  is  scarcely  soluble  in 
hydrochloric  acid,  but  partially  dissolves  in  excess  of  ammonium 
sulphide,  forming  a  brown  solution.  Three  other  sulphides  have 
been  obtained,  having  the  composition  Ni2S,  NiS2,  and  NigS^. 

2  X 


CHAPTER  XIV 

THE  TRANSITIONAL  ELEMENTS  OF  THE  SECOND 
AND  FOURTH  LONG  PERIOD 

Ruthenium,  Ru  =  101.7.  Rhodium,  Rh  =  io3.  Palladium,  106. 

Osmium,  05=191.  Iridium,  Ir=i93.  Platinum,  194.8. 

THESE  elements,  although  constituting  two  transitional  groups,  are  very  closely 
related  to  each  other.  In  nature  they  all  occur  associated  together  in  what  is 
commonly  known  as  platinum  ore,  and  they  are  on  this  account  usually  spoken 
of  as  the  platinum  metals. 

Platinum  ore,  or  native  platinum,  contains  all  these  elements  in  the  metallic 
state.  It  is  found  in  small  grains,  sometimes  in  nuggets,  in  alluvial  deposits  and 
river  sand,  principally  in  Brazil,  Borneo,  California,  Australia,  and  the  Urals. 
Native  platinum  contains  from  60  to  86  per  cent,  of  platinum,  the  remainder 
consisting  of  the  other  five  metals  of  the  group,  together  with  varying  quan- 
tities of  gold,  copper,  and  iron.  Amongst  the  grains  of  platinum  ore  there 
are  also  found  grains  which  consist  essentially  of  an  alloy  of  platinum  and 
iridium  (containing  from  30  to  75  per  cent,  of  iridium)  known  as  platin- 
iridium  :  and  also  particles  of  an  alloy  of  osmium  and  iridium  (called  osmiri- 
diurn],  which  contain  from  30  to  40  per  cent,  of  osmium,  as  well  as  small 
quantities  of  rhodium  and  ruthenium. 

They  are  all  white  lustrous  metals,  having  high  melting-points.  They  are 
unacted  upon  by  air  or  oxygen  at  ordinary  temperatures  ;  and,  with  the  excep- 
tion of  osmium  (which  burns  when  strongly  heated,  forming  the  tetroxide), 
they  are  scarcely  oxidised  by  oxygen  at  any  temperature. 

With  the  exception  of  palladium,  which  readily  dissolves  in  hot  nitric  acid, 
these  metals  are  unacted  upon  by  ordinary  acids.  Aqua  regia  converts 
osmium  into  the  tetroxide ;  it  dissolves  platinum  with  formation  of  the  tetra- 
chloride,  and  slowly  acts  upon  ruthenium,  but  is  without  action  upon 
rhodium  and  iridium. 

The  specific  gravities  of  the  metals  of  the  first  group,  although  very  close  to 
one  another,  are  widely  different  from  those  of  the  second  group ;  and  it  will 
be  seen  that  the  specific  gravities  fall,  with  increasing  atomic  weights,  thus — 

Ru,  sp.  gr.  =12.26.  Rh,  sp.  gr.  =  12.1.  Pd,  sp.  gr.  =11.4. 

Os,       ,,       =22.47.  Ir,         ,,       =22.38.  Pt,       ,,       —21.5. 

The  element  osmium  is  the  heaviest  known  substance. 

The  most  easily  fusible  of  these  metals  is  palladium,  which  melts  about  the 
temperature  of  wrought  iron.  The  melting-point  of  platinum  is  somewhat 
higher,  but  it  may  be  boiled  by  the  oxyhydrogen  flame.  Rhodium  and 

690 


Platinum  691 

iridium  come  next  in  order  of  fusibility,  the  latter  metal  being  just  fusible  by 
the  oxyhydrogen  flame,  while  ruthenium   has  a  still   higher   melting-point. 
Osmium  has  not  been  melted.     When  heated  to  the  melting-point  of  iridium, 
osmium  volatilises  ;  and  if  air  be  present,  it  burns. 
The  following  oxides  of  these  metals  are  known — 

Pd20 

RuO  OsO  RhO  PdO  PtO 

RuaOs         OsaOg          Rh2O3        Ir2O3 

RuO2          OsO2  RhO2         IrO2  PdO2  PtO2 

RuO4  OsO4 

Ruthenium,  osmium,  rhodium,  and  iridium  form  salts  corresponding  to  the 
sesquioxide,  such  as  ruthenious  chloride,  RuaClgj  rhodium  sulphate,  Rh2(SO4)3; 
iridious  chloride,  Ir2Cl6. 

With  the  exception  of  rhodium,  they  all  form  chlorides,  corresponding  to 
the  dioxides,  thus — ruthenic  chloride,  RuCl4;  iridic  chloride,  IrCl4 ;  platinic 
chloride,  PtCl4,  while  palladium  and  platinum  yield  palladia  and  plalinous 
compounds,  corresponding  to  their  monoxides. 

The  tetroxides  of  ruthenium  and  osmium  are  remarkable  in  melting  at 
an  extremely  low  temperature  (about  40°) ,  and  boiling  about  100°.  They 
yield  intensely  irritating  vapours,  which,  in  the  case  of  osmium  tetroxide, 
exerts  a  most  injurious  effect  upon  the  eyes,  and  is  extremely  poisonous. 
(Osmium  tetroxide  is  commonly  known  as  osmic  acid. )  Osmium  and  ruthenium 
also  exhibit  a  non-metallic  character  in  forming  compounds  derived  from  the 
unknown  ruthenic  and  osmic  trioxides,  such  as  potassium  ruthenate,  K2RuO4, 
and  potassium  osmate,  K2OsO4  (the  corresponding  ruthenic  and  osmic  acids 
are  unknown).  Ruthenium  also  forms  potassium  per-ruthenate,  KRuO4 
(analogous  to  permanganate),  although  the  corresponding  acid  and  peroxide, 
Ru2O7,  are  unknown.  The  most  important  of  these  elements  is  platinum. 


PLATINUM. 

Symbol,  Pt.     Atomic  weight  =  194. 8. 

In  order  to  separate  platinum  from  the  other  metals  with  which 
the  native  platinum  (see  page  690)  is  mixed,  the  ore  is  digested  in 
dilute  aqua  regia,  under  slightly  increased  pressure.  The  solution 
so  obtained  contains  the  higher  chlorides  of  platinum,  palladium, 
rhodium,  and  iridium  (for  although  in  the  pure  state  the  last  two 
named  metals  are  scarcely  attacked  by  aqua  regia,  when  alloyed  with 
much  platinum  they  dissolve).  The  solution  is  evaporated  to  dry- 
ness,  and  heated  to  125°,  whereby  the  palladium  and  rhodium  are 
obtained  in  the  form  of  their  lower  chlorides,  PdCl2  and  Rh2Cl(J 
(the  latter  of  which,  in  the  anhydrous  condition,  is  insoluble  in 
water).  The  residue  is  extracted  with  water,  and  to  the  clear  solu- 


692  Inorganic  Chemistry 

tion,  acidified  with  hydrochloric  acid,  ammonium  chloride  is  added. 
The  double  chloride  of  platinum  and  ammonium  (PtCl4,2NH4Cl), 
separates  out  as  yellow  crystals,  while  the  corresponding  iridium 
salt,  being  more  soluble,  remains  for  the  most  part  in  solution, 
and  may  be  obtained  by  concentrating  the  mother-liquor.  The 
ammonium  platinic  chloride,  on  being  ignited,  loses  ammonium 
chloride  and  chlorine,  leaving  the  metal  in  the  form  of  a  black 
spongy  mass  known  as  spongy  platinum,  which  is  then  melted  by 
means  of  the  oxyhydrogen  flame  in  a  lime  crucible.  The  platinum 
so  obtained  usually  contains  small  quantities  of  iridium  and  traces 
of  associated  metals. 

Pure  platinum  is  obtained  by  alloying  commercial  platinum 
with  pure  lead,  and  treating  the  alloy  first  with  nitric  acid,  which 
dissolves  any  copper  and  iron,  a  part  of  the  palladium  and  rhodium, 
and  most  of  the  lead  ;  and  then  with  dilute  aqua  regia,  which  dis- 
solves the  whole  of  the  platinum  and  the  remaining  lead,  with 
traces  of  rhodium.  From  this  solution  the  lead  is  precipitated  as 
sulphate,  and  the  platinum  is  then  precipitated  as  the  double 
chloride,  by  ammonium  chloride.  To  remove  traces  of  rhodium 
which  are  present,  the  dried  double  chloride  is  ignited  with 
hydrogen  potassium  sulphate,  whereby  the  rhodium  is  converted 
into  a  soluble  double  sulphate  of  rhodium  and  potassium,  while 
the  platinum  is  reduced  to  the  condition  of  the  spongy  metal. 

Properties. — Platinum  is  a  lustrous,  greyish-white,  malleable, 
and  ductile  metal.  At  a  red  heat  it  may  be  welded  with  great 
ease.  It  is  melted  by  the  oxyhydrogen  flame,  and  vessels  of 
platinum  are  readily  made  by  fusing  the  metal  together  in  this 
way.  Heated  platinum  absorbs  large  quantities  of  hydrogen 
(see  page  179) ;  and  when  the  metal  is  melted  in  the  oxyhydrogen 
flame,  it  exhibits  the  phenomenon  of  "spitting,"  when  it  again 
solidifies  (see  Silver,  page  562).  Platinum  does  not  combine  with 
oxygen  at  any  temperature,  neither  does  the  heated  metal  absorb 
this  gas  ;  but  it  has  the  property,  when  cold,  of  condensing  oxygen 
upon  its  surface.  A  piece  of  clean  platinum  foil  or  wire,  when 
introduced  into  a  mixture  of  oxygen,  and  a  readily  inflammable 
gas  or  vapour  (such  as  hydrogen,  ether,  alcohol,  &c.),  causes  their 
combination  ;  and  occasionally  the  metal  becomes  red  hot,  and 
ignites  the  mixture.  This  action  is  more  rapid  in  the  case  of 
platinum  sponge,  when  a  larger  surface  is  brought  into  play,  and 
a  fragment  of  this  material  introduced  into  a  detonating  mixture 
of  oxygen  and  hydrogen  at  once  determines  its  explosion. 


Platinum  Bichloride  693 

Platinum  is  not  acted  upon  by  either  nitric  or  hydrochloric  acid. 
It  is  oxidised  when  fused  with  caustic  alkalies,  or  with  potassium 
nitrate,  and  is  also  attacked  by  fused  alkaline  cyanides.  In  the 
form  of  sponge,  it  is  dissolved  by  boiling  potassium  cyanide,  with 
the  evolution  of  hydrogen  and  formation  of  a  double  cyanide. 

Platinum  readily  combines  with  phosphorus,  silicon,  and  carbon. 
The  carbide  of  platinum  is  formed  when  the  metal  is  continuously 
heated  by  a  smoky  flame,  or  one  in  which  combustion  is  incom- 
plete, hence  care  is  necessary  in  the  use  of  platinum  vessels. 

Platinum  Black  is  the  name  given  to  the  finely  divided  metal 
obtained  by  precipitating  platinum  from  its  solutions  by  reducing 
agents  or  by  metals.  It  is  a  soft,  black  powder,  which  is  capable 
of  absorbing,  or  condensing  upon  its  surface,  large  quantities  of 
oxygen.  It  therefore  acts  as  a  powerful  oxidising  agent. 

Platinum  Alloys.— Platinum  readily  alloys  with  many  metals  ; 
hence  compounds  of  easily  reducible  metals  should  not  be  heated 
in  vessels  of  platinum.  The  most  important  alloys  are  those  with 
iridium.  The  addition  of  2  per  cent  of  iridium  is  found  greatly  to 
increase  the  hardness  and  raise  the  melting-point  of  platinum. 
An  alloy  containing  10  per  cent,  of  iridium  resists  the  corrosive 
action  of  chemical  reagents  to  a  greater  extent  than  pure  platinum 
(see  Fluorine,  page  348). 

Oxides  of  Platinum.— Platinous  oxide,  PtO,  and  platinic  oxide, 
PtO2,  are  obtained  in  the  form  of  dark  grey  or  black  powders  by 
gently  heating  the  corresponding  hydroxides.  When  strongly 
heated  they  are  converted  into  the  metal. 

Platinous  Hydroxide,  Pt(HO)2,  is  obtained  by  the  action  of 
potassium  hydroxide  upon  platinum  dichloride.  It  is  a  black 
powder,  which  dissolves  in  the  halogen  acids,  yielding  plai'mous 
compounds. 

Platinic  Hydroxide,  Pt(HO)4,  is  prepared  by  adding  boiling 
potassium  hydroxide  to  a  solution  of  platinum  tetrachloride,  and 
treating  the  precipitate  with  acetic  acid  to  remove  the  potash. 
When  dried  it  forms  a  yellowish  powder,  which  is  soluble  in  acids 
to  form  platinic  salts.  Platinic  hydroxide  behaves  both  as  a  weak 
base  and  a  feeble  acid.  With  stronger  bases  it  forms  compounds 
known  as  platinates,  which  are  yellow  crystalline  salts.  The 
sodium  salt  has  the  composition  Na2O,3PtO2,6H2O. 

Platinum  Dichloride  (platinous  chloride),  PtCl2,  is  produced 
when  platinum  tetrachloride  is  heated  to  about  250°.  It  forms  a 
greenish  powder,  insoluble  in  water.  It  dissolves  in  hydrochloric 


694  Inorganic  Chemistry 

acid,  giving  a  reddish-brown  solution  which  is  believed  to  contain 
the  double  compound  PtCl2,2HCl,  or  H2PtCl4,  to  which  the  name 
chloro-platinous  acid  has  been  given.  The  compound  has  never 
been  isolated,  but  a  number  of  double  salts  of  platinous  chloride 
with  other  chlorides  are  known,  which  may  be  regarded  as 
derivatives  of  this  acid,  and  which  are  therefore  termed  chloro- 
platinites  ;  thus,  potassium  platinous  chloride,  2KCl,PtCl2,  or 
potassium  chloro-platinite,  K2PtCl4,  is  obtained  as  fine  red  crystals, 
by  adding  potassium  chloride  to  a  solution  of  platinous  chloride 
in  hydrochloric  acid.  This  salt  is  used  in  the  platinotype  photo- 
graphic process. 

Platinum  Tetraehloride  (platinic  chloride],  PtCl4,  is  obtained 
by  dissolving  the  metal  in  aqua  regia,  and  removing  the  excess 
of  the  acids  by  evaporating  to  dryness  and  gently  heating  the 
residue.  From  its  aqueous  solution,  the  salt  deposits  in  large 
red  crystals  having  the  composition  PtCl4,5H2O,  which  are  not 
deliquescent.  When  the  salt  is  crystallised  from  a  hydrochloric 
acid  solution,  or  when  the  aqua  regia  solution  is  evaporated  to 
expel  the  nitric  acid,  with  frequent  addition  of  hydrochloric  acid, 
the  double  compound  of  platinic  chloride  and  hydrochloric  acid  is 
formed,  PtCl4,2HCl,  which  is  deposited  as  reddish-brown  deli- 
quescent crystals,  with  6H2O.  To  this  substance  (which  is 
commonly  called  platinic  chloride),  the  name  chloro-platinic  acid 
has  been  given,  and  the  double  salts  of  platinic  chloride  and 
various  chlorides  are  regarded  as  salts  of  this  acid.  The  most 
important  of  these  chloro-platinates  are  those  of  the  alkali  metals, 
their  different  solubilities  being  made  the  basis  for  the  separation 
of  these  metals. 

Potassium  Chloro-platinate  (or  potassium  platinic  chloride), 
2KCl,PtCl4  or  K2PtCl6,  is  obtained  as  a  yellow  crystalline  pre- 
cipitate by  adding  potassium  chloride  to  platinic  chloride.  It  is 
soluble  in  loo  parts  of  water  at  the  ordinary  temperature  to  the 
extent  of  i.i  part,  and  at  100°,  5.18  parts.  It  is  insoluble  in  alcohol. 

The  rubidium  and  caesium  compounds  are  very  similar,  but  are 
still  less  soluble  in  water,  100  parts  of  water  at  20°  dissolving  0.141 
of  the  rubidium  and  0.07  of  the  caesium  salt. 

Ammonium  Chloro-platinate,  2NH4Cl,PtCl4,  closely  resembles 
the  potassium  salt,  being  slightly  less  soluble,  but  more  so  than 
the  rubidium  compound. 

Sodium  Chloro-platinate,  2NaCl,PtCl4,6H2O,  is  a  reddish* 
yellow  salt,  readily  soluble  in  both  water  and  alcohol. 


Platinamines  695 

Platino-eyanides. — Just  as  platinous  chloride  combines  with 
metallic  chlorides  to  form  chloro-platinites,  so  platinous  cyanide, 
Pt(CN)2,  unites  with  other  cyanides,  forming  similarly  constituted 
double  compounds,  known  as  platino-cyanides.* 

Potassium  platino-cyanide,  K2Pt(CN)4,  or  2KCN,Pt(CN)2,  is 
formed  when  spongy  platinum  is  dissolved  in  boiling  potassium 
cyanide.  The  platino-cyanides  may  be  regarded  as  the  salts  of 
platino-cyanic  acid,  H2Pt(CN)4.  Both  the  acid  and  the  salts  are 
characterised  by  the  wonderful  play  of  colours  they  exhibit  when 
viewed  in  different  lights,  and  by  forming  different  coloured 
crystals  with  varying  quantities  of  water  of  crystallisation  (see 
page  217). 

Sulphides  Of  Platinum.— Platinous  sulphide,  PtS,  and  platinic 
sulphide,  PtS2,  are  obtained  as  amorphous  black  powders  by  the 
action  of  sulphuretted  hydrogen  upon  the  respective  chlorides. 

Oxysalts  of  Platinum.— Few  well-defined  single  salts  of 
platinum  with  oxyacids  are  known.  This  element,  however, 
exhibits  a  great  tendency  to  form  complex  double  salts.  One  such 
series  of  compounds  is  seen  in  the  p  latino-nitrites ,  which  may  be 
regarded  as  the  salts  of  platino-nitrous  acid,  H2Pt(NO2)4. 

These  salts  are  remarkable,  in  that  the  platinum  they  contain 
cannot  be  detected  by  the  ordinary  tests  for  that  metal  ;  just  as 
the  iron  present  in  ferro-cyanides  is  not  detected  by  the  ordinary 
reagents  used  in  testing  for  that  metal. 

Ammoniacal  Platinum  Bases,  or  Platinamines. 

Like  cobalt,  platinum  forms  a  large  number  of  basic  compounds  with 
ammonia,  many  of  which  are  of  extremely  complex  composition.  The  first 
of  these  to  be  discovered  was  a  bright  green  salt,  obtained  by  the  action 
of  ammonia  upon  platinous  chloride,  having  the  composition  PtCl2,2NH3,  or 
Pt(NH3).,Clo,  and  known  as  the  green  salt  of  Magnus.  Many  of  the  platina- 
mines  exhibit  isomerism  ;  thus,  a  compound  known  as  the  chloride  of  Reisefs 
second  base  is  a  yellow  crystalline  salt  having  the  same  composition  as  Magnus's 
green  salt.  Twelve  distinct  series  of  ammoniacal  platinum  compounds  are 
known,  four  of  which  are  derived  from  platinous  and  the  remainder  from 
platimV  salts  ;  the  former  are  termed  platoso  ammonium  compounds,  while 
the  latter  are  distinguished  as  \.\\tp  latino  com  pounds,  f 

*  The  name  Cyano-platinites  might  with  advantage  be  applied  to  these 
compounds. 

t  For  detailed  descriptions  of  tfese  compounds,  the  student  is  referred  to 
larger  works  on  chemistry ;  and  on  the  constitution  of  these,  and  metallam- 
monium  compounds  generally,  the  article  by  Werner,  in  the  7.eitschrift  fur 
Anorganische  Chemie,  1893,  vol.  iii.  p.  267,  may  be  consulted. 


APPENDIX 


RADIUM,  AND  RADIOACTIVE  ELEMENTS 

As  far  back  as  the  year  1896,  Becquerel  discovered  that  the  element  uranium 
and  its  salts  possess  the  remarkable  property  of  emitting  rays  somewhat 
similar  in  character  to  the  now  familiar  Rontgen  or  "X"  rays;  resembling 
these  rays  in  their  penetrating  power,  their  photographic  action,  and  their 
action  upon  electrified  gases.  These  peculiar  rays  were  distinguished  from 
the  Rontgen  rays  by  being  called  the  "uranium,"  or  the  "  Becquerel"  rays. 
Somewhat  later  it  was  found  that  the  element  thorium  and  its  compounds 
were  likewise  possessed  of  the  property  of  emitting  rays,  which,  while  differing 
from  both  the  "X"  and  the  "uranium"  rays  in  some  respects,  closely 
resembled  them  in  others.  To  denote  this  property,  the  term  radioactivity 
has  been  coined,  and  substances  possessing  the  property  are  said  to  be 
radioactive  bodies. 

In  1898  it  was  announced  that  M.  and  Mme.  Curie  had  discovered  a  new 
radioactive  substance  contained  in  pitchblende,  a  mineral  consisting  essentially 
of  uranium  oxide.  From  researches  already  made,  it  had  been  shown  that 
the  radioactivity  of  uranium  compounds  is  roughly  proportional  to  the  amount 
of  the  metal  present,  but  it  was  found  that  in  the  case  of  certain  specimens  of 
pitchblende  this  was  not  the  case,  but  that  the  radioactivity  was  greatly  in 
excess  of  that  calculated  from  the  percentage  of  uranium  in  the  mineral.  This 
fact  suggested  the  presence  of  some  new  substance  of  superior  radioactivity  to 
that  possessed  by  uranium.  It  was  found  in  the  ordinary  process  of  separation 
of  the  metals  by  precipitation  from  an  acid  solution  by  sulphuretted  hydrogen, 
that  this  new  active  substance  was  thrown  down  along  with  the  sulphides,  and 
finally  was  separated  from  the  copper  and  arsenic,  &c. ,  and  remained  associ- 
ated with  the  bismuth.  No  isolation  of  the  new  substance  was  effected,  but 
from  its  greatly  superior  radioactivity  the  discoverers  concluded  that  there  was 
sufficient  evidence  of  the  presence  of  a  new  element  to  warrant  them  in  giving 
it  a  name.  They  therefore  called  it  polonium,  from  the  country  from  which 
the  pitchblende  was  obtained.*  (Compt.  rend.  127,  p.  175.) 

Following  up  their  investigations,  the  same  workers  very  shortly  afterwards 
discovered  in  the  same  mineral  another  radioactive  body  of  still  far  greater 
activity.  This  new  substance,  they  found,  is  not  precipitated  by  either 
sulphuretted  hydrogen,  ammonium  sulphide,  or  ammonia,  but  is  associated 

*  Although  the  name  polonium,  is  still  met  with  in  the  literature  of  the 
subject,  no  further  evidence  has  been  produced  in  proof  of  the  existence  of  a 
new  element  corresponding  to  the  name.  The  name  is  used  rather  to  denote 
the  radioactivity  which  appears  to  be  associated  with  the  element  bismuth. 

697 


698  Appendix 

with  and  accompanies  barium  in  the  various  'chemical  reactions  the  latter 
element  undergoes.  Thus,  when  barium  sulphate  or  carbonate  is  precipitated 
from  a  solution  of  the  chloride,  the  precipitated  barium  compound  is  accom- 
panied by  the  radioactive  material ;  or  when  the  chloride  itself  is  precipitated 
either  by  strong  hydrochloric  acid  or  by  alcohol,  the  "active"  substance  is 
thrown  down  along  with  it. 

By  the  careful  fractional  precipitation  of  the  chloride  with  alcohol  it  was 
found  possible  to  gradually  concentrate  the  radioactive  substance  in  the 
barium  chloride,  and  in  this  way  a  product  was  obtained  possessing  a  radio- 
activity 900  times  greater  than  that  of  uranium.  In  view  of  the  intensity 
of  its  "activity,"  the  discoverers  gave  the  name  radium  to  the  new  element 
which  they  believed  to  be  present,  although  in  almost  infinitely  minute 
quantities.  (Compt.  rend.  127,  p.  1215.) 

The  spectrum  exhibited  by  this  "active"  barium  chloride  also  confirmed 
the  presence  of  a  new  element,  for  besides  the  lines  belonging  to  barium  it 
contained  a  well-defined  line  which  had  never  previously  been  observed  in  the 
spectra  of  any  of  the  known  elements. 

Determinations  of  the  atomic  weight  of  the  metal  (barium)  in  the  speci- 
mens of  barium  chloride  which  contained  the  radioactive  element  to  an 
extent  sufficient  to  show  an  "activity"  900  times  greater  than  that  of 
uranium,  gave  values  practically  the  same  as  those  of  ordinary  barium,  namely 
137.4.  That  is  to  say,  the  actual  amount  of  radium  which  gave  rise  to  so 
high 'an  "activity"  in  the  barium  chloride  was  too  small  to  influence  the 
atomic  weight  determination.  When,  however,  the  concentration  of  the 
radium  chloride  in  the  barium  chloride  was  considerably  increased  by  a  con- 
tinuation of  the  fractionating  process,  the  atomic  weight  of  the  metal  was 
found  gradually  to  rise.  Thus,  when  the  intensity  of  the  radioactivity  reached 
3000  times  that  of  uranium,  the  atomic  weight  of  the  "  barium  "  rose  to  140 ; 
while  with  a  concentration  representing  a  radioactivity  7500  times  that  of 
uranium,  the  atomic  weight  of  the  metal  present  was  found  to  be  145.8.  From 
these  determinations  it  was  evident  that  radium  would  be  found  to  be  an 
element  of  very  high  atomic  weight,  and  in  the  course  of  time  when  it  became 
possible  to  obtain  small  quantities  of  radium  compounds— such  as  the  chloride 
and  bromide — in  a  state  of  comparative  purity,  this  was  found  to  be  the  case. 
The  latest  determinations  which  have  been  made  by  Mme.  Curie  and  others, 
have  assigned  the  number  226.4  as  the  atomic  weight  of  this  new  element. 
From  purely  spectroscopic  considerations,  however,  Runge  and  Precht  (Ast. 
Journ.,  April  1903)  calculate  the  atomic  weight  of  radium  to  be  258. 

The  element  radium  appears  to  resemble  barium  in  its  chemical  relations. 
Thus  the  sulphate  is  insoluble  in  water  and  in  acids  ;  the  carbonate  is  insoluble 
in  water,  and  the  chloride  is  precipitated  by  both  strong  hydrochloric  acid 
and  alcohol. 

As  seen  in  the  Bunsen  flame,  the  strongest  and  most  permanent  line  pro- 
duced by  radium  bromide  is  the  blue  line  4826. 

The  metal  itself  has  net  yet  been  isolated,*  and  in  view  of  the  extreme 


*  In  the  literature  of  the  subject  the  name  radium  is  constantly  employed 
when  in  reality  a  radium  salt  is  intended. 


Appendix  699 

minuteness  of  the  quantities  of  this  element  which  occur  in  the  mineral 
pitchblende  this  need  be  no  cause  for  surprise.  Not  only  is  the  amount  of 
radium  present  in  this  mineral  too  small  to  be  detected  by  any  chemical  test, 
but  the  spectroscope  itself  does  not  afford  a  sufficiently  delicate  means  for  its 
detection  ;  and  it  is  not  until  the  quantity  naturally  present  has  been  greatly 
concentrated  by  the  process  already  described,  that  the  characteristic  spectrum 
even  begins  to  make  its  appearance. 

The  chief  interest  attaching  to  this  new  element  is  associated  with  the 
strange  property  it  possesses  in  such  a  high  degree  of  emitting  "radiations." 
Radium  bromide,  for  example,  is  self-luminous  in  the  dark ;  the  rays  it  emits 
are  capable  of  acting  upon  a  photographic  plate,  much  as  the  Rontgen  rays 
affect  it.  They  cause  phosphorescence  upon  a  screen  of  barium  platino- 
cyanide,  and  produce  radiographic  effects  similar  to  those  given  by  the  "X" 
rays.  They  are  capable  of  penetrating  metals,  and  will  discharge  an  electro- 
scope not  only  through  considerable  intervals  of  space,  but  also  through 
screens  of  various  materials. 

Most  mysterious  of  all,  they  appear  to  possess  the  power  of  exciting  a 
temporary  radioactivity  in  other  substances  otherwise  inactive.  Thus,  if  a 
solution  of  a  radium  salt  and  some  distilled  water  are  placed  in  separate 
dishes  in  a  perfectly  closed  space,  radioactivity  is  communicated  to  the  water. 
The  water,  however,  gradually  loses  this  power  even  in  a  closed  space,  while 
it  rapidly  loses  it  if  exposed.  It  has  been  found  also  that  the  intensity  of  this 
"induced"  radioactivity  is  the  same  for  all  substances,  under  the  same  con- 
ditions, irrespective  of  their  chemical  nature. 

Concerning  the  nature  and  the  cause  of  the  radiations  emitted  by  radium 
and  the  other  two  well-defined  radioactive  elements  uranium  and  thorium,*  a 
large  amount  of  experimental  work  has  been  done,  and  much  speculation  put 
forward.  As  the  outcome  of  the  former  it  has  been  established  that  at  least 
four  distinct,  and  to  some  extent  separable  emissions,  may  go  to  make  up  what 
is  included  in  the  term  "  radiations."  These  are  distinguished  as  a,  /3,  and  y 
rays,  and  "  radioactive  emanation." 

i.  The  a  Rays, — These  rays  are  very  easily  absorbed  by  thin  layers  of 
matter.  Thus,  a  thickness  of  aluminium  0.0005  cm.  reduces  their  intensity 
to  one-half.  To  them  is  mainly  attributable  the  property  of  causing  the 
ionisation  of  a  gas,  whereby  its  electrical  conductivity  is  increased.  They 
are  deviated  by  a  very  strong  magnetic  field,  the  deviation  being  in  the 
opposite  direction  to  that  exhibited  by  "cathode"  rays.  These  a  rays  are 
not  waves  like  ordinary  light  rays,  but  consist  of  actual  matter,  which  is  being 
projected  at  an  enormous  velocity,  and  is  highly  charged  with  positive  electricity. 
They  are  described  as  a  "  flight  of  material  particles,"  having  a  mass  of  the 
same  order  as  the  atoms  of  hydrogen  ,f  and  travelling  with  a  velocity  about 
one-tenth  that  of  light. %  These  particles  carry  with  them  a  relatively  enor- 

*  Polonium,  and  the  still  more  recent  actinium,  are  at  present  too  undefined 
to  be  included  as  elements. 

rge  of  the  carrier  to  its  mass  is 

m 

%  That  is,  about  2.5  x  io9  cms.  per  sec.  (Rutherford  and  Soddy,  Phil.  Mag., 
Feb.  1903). 


700  Appendix 


mous  amount  of  energy,  each  particle  apparently  having  sufficient  energy 
associated  with  it  to  excite  phosphorescence  visible  to  the  eye.  Thus,  Crookes 
has  shown  that  when  a  fragment  of  solid  radium  nitrate  is  brought  near  to  a 
screen  of  "  Sidot's hexagonal  blende"  (zinc  sulphide),  and  the  phosphorescent 
surface  of  the  screen  is  examined  with  a  pocket  lens,  it  is  seen  to  be  dotted  all 
over  with  brilliant  specks  of  green  light.  In  proportion  as  the  radium  salt 
is  brought  closer  to  the  screen,  these  flashes  or  scintillations  become  more 
brilliant  and  more  numerous,  following  each  other  with  such  rapidity  that  the 
surface  presents  the  appearance  of  a  "  turbulent  luminous  sea." 

"  It  seems  probable  that  we  are  here  actually  witnessing  the  bom- 
bardment of  the  screen  by  the  electrons  hurled  off  by  the  radium" 
(Crookes). 

2.  The  /3  Rays. — These  rays  are  readily  deviated  by  the  magnetic  field; 
and  also  differ  from  the  a  rays  in  their  greater  penetrating  powers.     Thus, 
while  the  intensity  of  the  latter  is  reduced  to  one-half  by  passing  through 
0.0005  crn-  °f  aluminium,  the   ^3  rays  are  able   to  traverse  a  thickness  of 
0.05  cm.  of  this  metal  before  their  intensity  is  halved.    A  sheet  of  mica  o.oi  cm. 
thick  will  completely  absorb  all  the  a  rays,  while  it   transmits   the  j3  and 
also  the  7  rays  without  appreciable  diminution.      The   /3   rays,    like   the   a 
rays,  also  consist  of  projected  particles  with  a  high  velocity,  but  in  this  case 
they  carry  a  negative  electric  charge,  and  their  mass  is  believed  to  be  greatly 
less  than  that  of  the  particles  constituting  the  a  rays,  namely,  about  the  ^^nny  °f 
that  of  the  hydrogen  atom  (Rutherford  and  Soddy,  Phil.  Mag.,  May  1903). 
/3  rays  are  similar  in  all  respects  to  the  "cathode  "  rays  emitted  from  a  vacuum 
tube,  except  that  the  velocity  of  the  particles  is  greater  and  consequently  they 
are  more  penetrative.     Their  velocity  is  estimated  to  be  between  2  x  io10  and 
3  x  io10. 

3.  The  ^  Rays.  —  These  are  non-deviable  by  the  magnetic  field,  and  closely 
resemble  the  Rontgen  or  "X''rays.     They  are  far  more  penetrating  than 
either  the  a  or  f3  rays,  being  capable  of  penetrating  a  thickness  of  8.0  cms. 
of  aluminium  before  their  intensity  is  reduced  to  one-half.      These  rays  are 
believed  to  be  a  wave  motion,  and  not  to  consist  of  projected  particles  of 
matter. 

4.  "  Radioactive  Emanation." — The  elements  thorium  and  radium  *  possess 
the   property   of  emitting  something   which    has    the    power    of   imparting 
radioactivity   to    any  substance    in    their  immediate  neighbourhood.       The 
radioactivity  thus  imparted   or    excited    is  only  of  a   temporary  character, 
its  intensity  diminishing   and   dying  away  when  the   substance  is  removed 
from  the   influence   of  the  original  radioactive     body.      Experiments   seem 
to  prove   that   these  effects   are   not   produced   by  any  of  the  rays  already 
described,    but   are    due    to   some    other    distinct   emission,    and    the    term 
"radioactive  emanation,"   or  shortly    "emanation,"  has    been   adopted   to 
denote  this. 

The  radioactivity  which  is  thus  imparted  to  substances  in  the  proximity 
of  these  radioactive  elements  (usually  spoken  of  as  excited  radioactivity)  is 
believed  to  be  caused  by  the  deposition  upon  their  surface  of  radioactive 
matter,  which  is  transmitted  by  positively  charged  carriers  ;  while  the  radio- 

*  Uranium  appears  not  to  share  this  property. 


Appendix  701 

activity  of  the  "  emanation  "  itself  is  believed  to  be  due  to  the  emission  from 
it  of  a  rays  only.  When  a  small  quantity  of  thorium  oxide*  is  placed  in  a 
tube  (the  oxide  being  enveloped  in  material  capable  of  intercepting  the 
ordinary  radiations)  and  a  stream  of  air  is  passed  over  it,  the  air  is  found 
to  carry  with  it  the  "emanation"  which  the  thorium  oxide  gives  out;  and 
the  issuing  stream  of  air,  even  after  being  conveyed  through  many  feet  of 
tube,  is  capable  of  discharging  an  electroscope.  In  the  case  of  radium 
compounds  the  amount  of  this  "emanation"  was  found  to  be  comparatively 
small  when  the  radium  compound  is  employed  in  the  solid  state,  but  when 
the  radium  salt  is  dissolved  in  water,  the  "emanation"  appears  to  be  given 
off  in  a  sudden  rush,  as  it  were,  and  the  solution  continues  to  emit  this 
"emanation"  in  amount  many  hundred  times  as  great  as  was  produced  by 
the  solid  salt.  A  similar  enormous  increase  also  takes  place  when  the  radium 
compound  is  heated.  These  observations  have  led  to  the  belief  that  the 
"  emanation  "  is  actually  occluded  by  the  solid  compound,  f 

In  many  other  respects  this  "emanation"  behaves  like  an  inert  gas.  Thus 
if  the  stream  of  air  carrying  the  "emanation"  is  passed  through  a  tube 
plugged  with  cotton  wool,  nothing  is  arrested  or  filtered  out  by  the  wool  and 
the  radioactivity  of  the  air  as  it  issues  is  not  diminished.  Neither  is  it  affected 
by -the  air  being  bubbled  through  strong  sulphuric  acid,  or  passed  through  a 
red-hot  platinum  tube.  When  air  conveying  "  emanation  "  is  slowly  passed 
through  a  LJ-tube  cooled  by  liquid  air,  the  "emanation"  is  completely  con- 
densed, and  the  air  which  passes  out  is  entirely  free  from  all  trace  of  this 
substance.  If  a  glass  tube  is  employed,  and  the  air  current  is  sufficiently 
slow,  the  progress  of  the  condensation  can  be  traced  by  the  fluorescent 
appearance  of  the  glass,  showing  that  the  condensation  has  all  taken  place 
upon  the  first  portions  of  the  tube  traversed  by  the  stream  of  air.  If  now  the 
tube  is  closed  at  both  ends  and  the  temperature  allowed  to  rise  above  a 
certain  point,  the  condensed  "emanation"  appears  to  vaporise  again,  and 
the  fluorescence  extends  throughout  the  entire  length  of  the  tube.  The 
volatilisation  point  of  radium  emanation  appears  to  be  about  -  150°,  while 
that  of  the  thorium  emanation  is  given  as  about  — 120°  (Phil.  Mag. ,  1903, 

P-  575)- 

In  the  case  of  thorium,  the  "emanation"  loses  its  radioactivity,  or  decays, 
much  more  rapidly  than  the  radium  emanation.  Thus,  while  the  activity 
of  thorium  emanation  falls  to  half  its  intensity  in  the  space  of  one  minute, 
the  intensity  of  the  radium  emanation  only  sinks  to  half  its  value  in  the  space 
of  four  days,  while  still  retaining  sufficient  activity  to  be  detected  after  the 
lapse  of  one  month.  This  rate  of  decay  of  the  radioactivity  of  the 
"emanation"  is  the  same  even  at  the  low  temperature  of  liquid  air,  and  it 
is  considered  probable  that  the  marked  difference  in  the  rates  of  decay  of  the 
"emanation"  from  thorium  and  radium  may  account  for  the  difference 
observed  in  their  vaporisation  temperatures. 

It  was  at  one  time  supposed  that  the  radioactivity  of  these  radioactive 
elements  was  not  a  property  intrinsic  to  the  elements  themselves,  but  was  due 

*  Most  of  the  earlier  work  by  Rutherford  and  Soddy  (Phil.  Mag. ,  1902)  in 
this  connection  was  done  with  thorium  compounds, 
f  Rutherford  and  Soddy,  Phil.  Mag. ,  1903,  p.  449. 


jQ2  Appendix 

to  the  presence  in  small  and  varying  quantity  of  some  unknown  substance. 
Crookes  found  (Proc.  Royal  Soc.,  1900)  that  by  processes  of  a  purely  chemical 
nature  he  was  able  to  separate  from  uranium  nitrate  small  quantities  of 
material  which  seemed  to  possess  all  the  radioactivity,  leaving  the  bulk 
of  the  uranium  compound  inactive.  He  applied  the  name  Uranium  X  to  this 
"  unknown  uranium."  Similarly  in  the  case  of  thorium  ;  when  the  hydroxide 
was  precipitated  by  ammonia,  and  the  filtrate  (which  chemically  should  contain 
no  thorium)  was  evaporated  to  dryness  and  ignited  to  expel  ammonia  salts, 
minute  residues  were  obtained  which  were  many  hundred  times  more  active 
than  an  equal  weight  of  thorium  oxide  (Rutherford  and  Soddy,  Phil.  Mag., 
September  1902).  The  precipitated  hydroxide,  although  not  entirely  robbed 
of  radioactivity,  was  found  to  have  its  activity  greatly  reduced.  This  sup- 
posed "active  "  constituent  was  therefore  called  Th  X.  Later  investigations, 
however,  revealed  the  remarkable  fact  that  the  thorium  compound  which  had 
thus  been  partially  deprived  of  its  radioactivity  gradually  regained  it  when 
left  to  itself;  while  the  separated  Th  X  gradually  lost  it.  Moreover,  it  was 
found  that  the  two  processes  went  on  exactly  at  the  same  rate,  that  the  rate 
of  decay  of  the  activity  of  Th  X  was  the  same  as  the  rate  of  recovery  of  activity 
of  original  thorium  compound.  From  this  it  would  appear  that  two  opposing 
processes  are  simultaneously  going  forward  in  a  radioactive  substance,  namely, 
the  continual  production  of  fresh  radioactive  material  and  the  constant  decay 
of  the  radiating  power  of  the  active  material.  In  other  words,  what  may 
be  called  the  normal  radioactivity  is  a  condition  of  equilibrium,  where  the 
rate  of  increase  of  activity  due  to  the  production  of  fresh  active  material 
balances  the  rate  of  the  decay  of  the  activity  in  the  radioactive  material 
already  formed. 

The  views  now  generally  held  are  that  the  phenomena  of  radioactivity  are 
due  to  atomic  changes,  but  changes  of  a  character  altogether  different  from 
any  that  have  previously  been  dealt  with  in  chemistry.  It  is  believed  that  the 
atoms  of  these  radioactive  elements  (which,  it  will  be  noted,  are  possessed  of 
the  highest  atomic  weights  of  all  the  elements)  are  undergoing  a  process  of 
disintegration  or  degradation  :  that  in  the  course  of  their  movements,  owing 
to  some  combination  of  conditions  about  which  at  present  we  know  nothing, 
the  kinetic  energy  of  some  of  the  atoms  reaches  a  point  beyond  which  the 
stability  of  the  atom  is  no  longer  possible.  Under  these  circumstances  the 
atom  breaks  up,  throwing  off  some  matter  from  itself,  and  assumes  a  more 
stable  configuration.  The  particles  or  fragments  of  the  original  atoms  them- 
selves undergo  further  changes,  giving  off  other  particles,  thus  giving  rise  to 
the  various  phenomena  of  radioactivity. 

In  the  case  of  radium  there  seems  to  be  indubitable  evidence  (based  on  the 
work  of  Ramsay  and  others)  that  one  of  the  final  products  of  the  radioactive 
change  is  the  element  helium.  A  minute  quantity  of  radium  emanation 
enclosed  in  a  Pliicker  vacuum  tube  was  found  after  the  lapse  of  a  few  days  to 
give  the  characteristic  spectrum  of  helium. 

It  has  been  further  suggested  (Boltwood,  Phil.  Mag.,  1905)  that  radium 
itself  is  a  product  of  the  radioactivity  of  uranium  ;  not  necessarily  a.  first  pro- 
duct, but  probably  through  one  or  more  intermediate  stages. 

If  these  are  the  true  interpretation  of  the  phenomena  of  radioactivity  we  are 
undoubtedly  face  to  face  with  an  actual  instance  of  the  "  transmutation  of  the 


Appendix  703 

elements,"  which  has  hitherto  been  regarded  only  as  an  idle  dream  of  the 
alchemist.  It  may,  indeed,  be  that  in  these  radioactive  processes  we  have  as 
it  were  a  peep  into  the  unknown  region  of  the  "  evolution  of  the  elements."  * 

The  energy  which  is  liberated  during  this  process  of  atomic  disintegration  is 
enormous,  taking  into  account  the  minute  quantities  of  matter  concerned. 
M.  and  Mme.  Curie  have  shown  that  a  sample  of  a  radium  salt  gave  out 
energy  sufficient  to  melt  half  its  own  weight  of  ice  per  hour.  This  energy, 
which  is  stored  up  in  the  atoms  of  these  elements,  the  "  internal  energy  of  the 
chemical  atom,"  as  it  has  been  termed,  and  which  is  set  nee  during  radioactive 
change,  is  of  an  entirely  different  order  of  magnitude  from  that  which  is  dis- 
engaged during  any  processes  of  ordinary  chemical  change.  It  has  been 
calculated,  indeed,  that  the  energy  of  radioactive  change  is  many  thousand 
times,  or  even  a  million  times,  as  great  as  that  of  any  known  chemical 
change,  when  equal  weights  of  matter  are  concerned.  Ramsay  has  recently 
shown  f  that  this  energy  is  capable  of  bringing  about  certain  ordinary  chemical 
changes.  Thus  radium*  "  emanation  "  is  able  both  to  decompose  water  into 
its  elements,  and  also  to  cause  the  recombination  of  oxygen  and  hydrogen. 
It  is  always  found,  however,  that  the  mixed  gases  resulting  from  the  action  of 
the  "emanation"  upon  water,  contain  a  slight  excess  of  hydrogen,  the  exact 
reason  of  which  is  at  present  unknown. 

The  idea  of  an  atom  as  a  system,  and',  moreover,  one  capable  of  under- 
going changes  into  simpler  systems,  is  a  view  which,  to  the  chemist,  may  at 
first  seem  strangely  heterodox,  and  one  altogether  opposed  to  fundamental 
doctrines  of  chemistry.  In  reality,  however,  this  new  view  as  to  the  con- 
stitution of  an  atom  does  not  touch  the  question  of  the  indivisibility  of  the 
atom  in  the  purely  chemical  sense.  From  this  point  of  view  the  chemical 
atom  still  retains  its  position  as  the  lowest  stage  in  the  complexity  of  matter, 
and  may  still  be  defined  as  the  smallest  particle  of  matter  which  can  take 
fart  in  a  chemical  change.  The  chemical  atoms  of  these  radioactive  elements 
are  not  divisible  into  what  may  be  called  "  chemical  fragments."  If  the  atom 
is  a  system,  then  in  all  chemical  reactions  and  changes  the  system  in  its 
entirety  takes  part.  When  it  is  borne  in  mind  that  the  weight  of  matter 
which  the  atom,  regarded  as  a  changing  system,  throws  off  in  the  form  of 
"radiations,"  "emanation,"  or  "electrons"  is  so  infinitely  minute,  that  it  has 
been  estimated  that  it  would  require  many  hundreds,  if  not  thousands  of 
years  before  enough  of  it  could  be  collected  to  be  detected  by. the  most  deli- 
cate balance,  it  will  be  evident  that  we  are  dealing  with  phenomena  of  a  totally 
different  order  from  those  in  which  the  relative  weights* of  matter  entering  into 
chemical  combination  are  concerned. 

*  A  comprehensive  theory  of  the  evolution  and  devolution  of  the  elements, 
by  Jessop,  is  to  be  found  in  the  Phil.  Mag.,  Jan.  1908. 
f  Jour.  Chem.  Soc. ,  1907. 


INDEX 


ABSOLUTE  boiling-point,  79 

,,       temperature,  70 
Absorptiometer  (Bunsen),  144 
Absorption  of  gases  by  charcoal, 
Acetylene,  317 
Acetylide  of  copper,  318 
Acid,  ammon-sulphonic,  282 
,,  .    antimonic,  497 
,,       arsenic,  488 
,,      arsenious,  487 
,,      boric,  609 

bromic,  383 

,,      carbamic,  311,  444 
,,      carbonic,  305 
,,       chloric,  374 
,,      chloro-auric,  569 
,,       chlorochromic,  664 
,,       chloroplatinic,  694 
,,      chloroplatinous,  694 
,,       chlorosulphuric,  439 
,,       chlorosulphonic,  439 
,,       chromic,  662 
,,       dithionic,  437 
,,       hydrazoic,  279 
.,       hydriodic,  389 
,,      hydrobromic,  381 
,,      hydrochloric,  363 
„       hydrofluoboric,  612 
„       hydrofluoric,  350 
,,      hydrofluosilicic,  632 
,,       hydrosulphurous,  423 
,,       hypobromous,  383 
hypochlorous,  373 
,,       hypoiodous,  395 
,,       hyponitrous,  250 
,,       hypophosphorous,  472 
,,       hyposulphuric,  437 
,,      hyposulphurous,  423 
„      iodic,  392 


292 


705 


Acid,  manganic,  660 

,,  metaboric,  610 

metantimonic,  497 

,,  metaphosphoric,  476 

,,  metarsenic,  488 

,,  metasilicic,  635 

,,  metastannic,  640 

,,  metatungstic,  665 

,,  metavanadic,  656 

,,  molybdic,  664 

,,  muriatic,  371 

,,  nitric,  234 

,,  nitrosulphuric,  426 

,,  nitrous,  244 

,,  Nordhausen  sulphuric,  434 

,,  ortho-antimonic,  496 

, ,  ortho-arsenic,  488 

,,  ortho-arsenious,  487 

,,  orthoboric,  609 

,,  orthophosphoric,  474 

,,  orthosilicic,  635 

,,  osmic,  691 

,,  oxymuriatic,  352 

,,  pentathionic,  438 

,,  perchloric,  375 

,,  perchromic,  660 

periodic,  393 

,,  permanganic,  669 

persulphuric,  424 

,,  phosphomolybdic,  665 

,,  phosphoric,  474 

,,  phosphoric  (glacial),  476 

,,  phosphorous,  473 

,,  pyro-antimonic,  496 

,,  pyro-arsenic,  488 

,,  pyro-arsenious,  487 

,,  pyroboric,  610 

,,  pyrophosphamic,  478 

,,  pyrophosphodiamic,  478 
2  Y 


706 


Index 


Acid,  pyrophosphoric,  475 

,,       pyrophosphotriamic,  478 
pyrosulphuric,  434 

,,       pyrovanadic,  656 

,,       selenic,  447 

,,      selenious,  447 

,,       silicic,  635 

,,       stannic,  640 

,,       sulphovinic,  315 

,,       sulphuric,  425 

,,       sulphurous,  421 

,,       telluric,  449 

,,      tellurous,  449 

,,       tetrathionic,  438 

,,       thiocarbamic,  444 

,,       thiocarbonic,  443 

,,       thiosulphuric,  435 

,,       trithionic,  437 

,,       tungstic,  664 
Acid-forming  oxides,  17 
Acids,  dibasic,  19 

,,      mono-,  tetra-,  and  tribasic,  19 
Active  mass,  94 
Affinities,  61 

Affinity,  chemical,  10,  61 
After-damp,  298 
Air-liquefiers,  77 
Alabaster,  581 
Algin,  386 
Alkali  manufacture,  534 

metals,  505 
Alkali- waste,  400 
Alkaline  earths,  571 
Allotropy,  194 
Aludels,  386,  598 
Alum,  620 

,,       burnt,  622 

,,       meal,  621 

,,       shale,  621 
stone,  621 
Alumina,  617 
Aluminates,  618 
Aluminite,  619 
Aluminium,  614 

,,       alloys,  617 

,,       bronze,  553,  617 

,,       chloride,  623 

,,       fluoride,  622 

,,       hydroxides,  618 

,,       sodium  chloride,  623 


Aluminium  sulphate,  619 

,,       sulphide,  617,  623 
Alums,  620 
Alunite,  621 

Amalgamation  process  (silver),  559 
Amalgams,  600 
American  pot-ashes,  521 
Amethyst,  617 
Ammonia,  272 

,,       solubility  of,  in  water,  275 
Ammonia-soda  process,  538 
Ammoniacal  cobalt  compounds,  685 

,,       liquor,  319 

,,       mercury  compounds,  604 

,,       platinum  compounds,  695 
Ammonium,  545 

,,       alum,  620 

amalgam,  545 

,,       borofluoride,  612 

,,       carbamate,  311,  547 

,,       carbonate,  547 

,,       chloride,  546 

,,  ,,        dissociation  of,  89 

,,       chloroplatinate,  694 

,,       chromate,  230 

,,       cyanate,  14,  24 

,,       ferrous  sulphate,  680 

,,       hydrazoate,  280 

,,       hypoiodite,  283 
iron  alum,  620 

,,       magnesium  arsenate,  488 

,,       magnesium  phosphate,  475 

,,       manganous  chloride,  668 
meta-thio-arsenate,  490 
metavanadate,  656 

,,       molybdate,  665 
nitrate,  248 

,,       nitrite,  230 

phosphomolybdate,  477,  665 

,,       plumbic  chloride,  650 

,,       pyro-arsenite,  487 

,,       pyro-thio-arsenite,  490 

,,       salts,  545 

sesquicarbonate,  548 

,,       sodium  phosphate,  475 
stannic  chloride,  642 

,,       sulphate,  546 

,,  thiocyanate,  548 
Ammon-sulphonates,  282 
Amorphous  silicon,  629 


Index 


707 


Analysis,  13 
Anastase,  627 
Anglesite,  643 
Anhydrides,  17 
Anhydrite,  577 
Animal  charcoal,  290 
Anions,  99,  105 
Anodes,  97 
Anthracite,  293 
Antimonates,  497 
Antimonious  oxide,  496 
Antimony,  491 

,,       amorphous,  492 

,,       blende,  491 

,,       bloom,  491 

,,       compounds    with     halogens, 

494 

,,       chlorides,  495 

,,       hydride,  493 

,,      ochre,  491 

,,       oxides  and  oxyacids,  496 

,,       oxychlorides,  495 

,,       sulphides,  498 

,,      sulpho-trichloride,  496 

,,       tetroxide,  497 

,,       trioxide,  496 
Apatite,  347,  582 
Apollinaris  water,  220,  305 
Aquafortis,  234 
Aqua  regia,  241 

Aqueous  vapour  (atmospheric),  257 
Argentic     compounds     (see     Silver), 

558 

Argentiferous  lead,  560 
Argentite,  558 
Argon,  265 

Argon  group  of  gases,  the,  263 
Arragonite,  582 
Arsenates,  488 
Arsenic,  478 

,,      allotropic    modifications    of, 
480 

,,       chlorhydroxide,  484 

,,       chloride,  483 

compounds     with     halogens, 
483 

,,       fluoride,  483 

,,       hydride,  480 

,,       oxides  and  oxyacids,  484 

,,       pentoxide,  487 

„       sulphides,  489 


Arsenical  iron,  479 
pyrites,  479 
Arsenious  bromide,  484 

,,       iodide,  484 

,,       oxide,  484 
Arsenites,  487 

Arsenuretted  hydrogen,  480 
Arsine,  480 
Asbestos,  573 
Asymmetric  system,  162 
Atacamite,  555 
Atmolysis,  84 
Atmosphere,  252 

,,       composition  of,  256 

, ,      height  of,  263 

,,       suspended  impurities  in,  261 
Atmospheric  ammonia,  258 

,,       aqueous  vapour,  257 

,,       carbon  dioxide,  257 

gases  mechanically  mixed ,260 

,,  ,,     argon  group  of,  263 

,,       hydrogen,  260 

,,       nitric  acid,  258 

,,       ozone,  259 
Atomic  electric  charge,  104 

„       heat,  46 
theory,  25 

,,       volumes,  44,  119 

,,       weight,  definitions  of,  37,  44 

,,       weight,  determination  of,  by 

chemical  methods,  36,  58 

weight,  determination  of,  by 

means  of  isomorphism,  51 

,,       weight,  determination  of,  by 
means  of  specific  heat,  45 
weight,  determination  of.from 
volumetric  relations,  38 

,,       weights,  list  of,  22 

,,         international,  22,  38 
Atoms,  4 
Aurates,  569 
Auric  chloride,  569 

,,       oxide,  569 
Auro-auric  sulphide,  569 
Aurous  iodide,  569 
Autogenous  soldering,  430 
Avogadro's  hypothesis,  40 
Axes  of  symmetry,  160 
Azoimide,  279 
Azote,  229 
Azurite,  550 


Index 


BALANCED  actions,  88 
Balling  furnace,  536 
Barium,  586 

,,      amalgam,  586 

,,       bromate,  384 

,,       carbonate,  586 

,,      chlorate,  374 

,,       chloride,  588 

,,       dioxide,  184,  587 

,,       dithionate,  437 

,,       hydroxide,  587 

,,       hypophosphite,  472 

,,       iodate,  586 

,,       monoxide,  586 

,,       nitrate,  589 

,,       oxides,  586 

,,       peroxide,  587 

,,      sulphate,  588 

,,       sulphide,  589 

,,       tetrathionate,  438 

,,       thiosulphate,  438 
Baryta,  586 

,,      water,  587 
Barytocalcite,  586 
Base,  1 8 
Basic  oxides,  17 

,,       salts,  20 

Basicity  of  acids,  the,  18,  473 
Battery,  galvanic,  96 
Bauxite,  614 
Beryl,  572 
Berylla,  572 
Beryllium,  572 

,,       aluminate,  618 

,,       compounds,  572 

,,      specific  heat  of,  48 

,,       Bessemer  process  (steel),  675 
Binary  compounds,  15 
Bismuth,  500 

,,       alloys,  500 

, ,       carbonate,  502 

,,       compounds     with     halogens, 

501 

,,       dichloride,  501 
,,      dioxide,  502 
,,       glance,  500 
nitrate,  503 
,,       nitrate,  basic,  503 
,,      ochre,  500 
,,       oxides,  501 
,,       oxychloride,  503 


Bismuth  pentoxide,  504 

,,       tetroxide,  503 

,,       tribromide,  501 

,,       trichloride,  501 

,,       tri-iodide,  501 

,,       trioxide,  502 

,,       trisulphide,  504 
Bismuthic  oxide,  501 
Bismuthous  oxide,  501 
Bisulphate  of  soda,  422 
Bittern,  531 
Bituminous  coal,  293 
Black  ash,  composition  of,  537 
furnace,  535 

,,       revolving  furnace,  536 
Black-band,  672 
Black-jack,  591 
Blacklead,  288 
Blast-furnace,  673 
Bleaching-powder,  187,  373,  580 
Blister  copper,  551 

steel  675 
Blue  vitriol,  556 
Boiling-point,  absolute,  79 

,,     .  definition  of,  128 

, ,       molecular  elevation  of,  134 
Boiling-points,  129 

,,       effect  of  pressure  upon,  129 

,,      effect  of  dissolved  substances 
upon,  134 

,,       of  saturated  saline  solutions, 

J33 

Bolognian  phosphorus,  584 
Bone  ash,  452 

,,       black,  290 
Bones,  composition  of,  291 
Boracite,  607 
Borate  spar,  607 
Borates,  610 
Borax,  610 
Borofluorides,  612 
Boron,  607 

,,       hydride,  613 

,,       nitride,  613 

,,      sulphide,  613 

,,       trichloride,  612 

,,•     trifluoride,  611 

,,       trioxide,  608 
Boronatrocalcite,  607 
Bort,  285 
Boyle,  law  of,  71 


Index 


709 


Brass,  553 
Braunite,  666 

Erin's  process  (oxygen),  184 
Britannia  metal,  492,  639 
Brb'ggerite,  264 
Bromates,  384 
Bromides,  382 
Bromine,  377 

,,       electrolytic      manufacturing 
process,  380 

,,       hydrate,  380 

,,       monochloride,  396 

,,       oxyacids,  383 

,,       water,  380 
Bromous  acid,  383 
Bronze,  639 
Brookite,  627 
Brown  haematite,  672 
Brown  iron  ore,  672 
Brucite,  574 
Bunsen  flame,  the,  341 

,,       non-luminosity  of,  342 

,,       temperature  of,  343 
Burnt  alum,  622 

CADMIUM,  596 

,,       chloride,  597 
,,       oxide,  597 
,,       sulphide,  597 
Caesium,  505,  544 

,,       spectrum  of,  510 
Cailletet's  apparatus,  75 
Calamine,  591 
Calcined  magnesia,  574 
Calcite,  576 
Calcium,  577 

bicarbonate,  221,  311,  582 

borate,  611 

borofluoride,  612 

carbide,  319,  455 

carbonate,  582 

chlorate,  517 

chloride,  579 

chloro-hypochlorite,  580 

dioxide,  579 

fluoride,  347 

hydride,  578 

hydroxide,  578 

hypochlorite,  517,  580 

manganite,  359,  668 

oxides,  578 


Calcium  phosphate,  452,  582 

,,       phosphide,  460 

,,       sulphate,  581 

,,       sulphide,  400,  410,  583 
Calc-spar,  577 
Caliche,  387 
Calomel,  602 
Calorie,  165,  326 
Calx,  252,  321 
Candle  flame,  335 
Canton's  phosphorus,  584 
Capillary  pyrites,  689 
Carat,  definition  of,  569 
Carbides,  295 

,,       of  iron,  513 
Carbon,  285,  627 

,,       compounds,  295 

,,       dioxide,  300 

,,  ,,       atmospheric,  257 

,t  i,       composition  of,  307 

>i  ii       solid,  307 

,,       disulphide,  441 

,,       hydrogen,  compounds  of,  312 

,,       monoxide,  296 

,,       oxides  of,  296 

,,      specific  heat  of,  47 

,,      suboxide,  311 
Carbonado,  285 
Carbonates,  310 
Carbonyl  chloride,  299 
Carbonyls,  metallic,  299 
Carborundum,  630 
Car  boxy-haemoglobin,  298 
Carnallite,  512,  579 
Caro's  acid,  425 
Carre"  s  freezing-machine,  132 
Cassiterite,  637 
Cast  iron,  674 
Catalysis,  183 

Catalytic  action,  12,  183,  354 
Cathodes,  97 
Cations,  99,  105 
Caustic  potash,  515 

,,      soda,  530 
Celestine,  584 
Cellulose,  240 

Cementation  process  (steel),  675 
Centres  of  symmetry,  160 
Cerite,  627 
Cerium,  627 
Cerussite,  643 


710  Index 

Chalcedony,  628 

Chalk,  582 

Chalybeate  waters,  220 

Chamber  acid,  429 

Chamber  crystals,  426 

Chance's  process,  411 

Change  of  volume  on  solidification, 

137 
Charcoal,  290 

,,       absorption  of  gases  by,  292 
,,       animal,  290 
,,       specific  heat  of,  47 
Charles'  law,  69 
Chemical  action,  n 
,,       affinity,  10 
,,       combination,  laws  of,  25 
,,       equations,  23 
,,       formulae,  23 
,,       modes  of,  13 
,,       nomenclature,  15 
,,       notation,  quantitative,  53 
,,      reactions,  23 
,,       symbols,  22 
Chili  saltpetre,  541 
Chlorates,  375 
Chloride  ions,  104 
Chloride  of  lime,  580 
Chlorine,  352 

,,       heptoxide,  373 
,,       hydrate,  362 
,,       liquefaction  of,  73,  362 
,,       liquid,  362 

, ,       manufacturing  processes,  354 
,,       monoxide,  371 
,,       oxides  and  oxyacids,  371 
>,       peroxide,  372 
,,      water,  361 
Chloro-aurates,  569 
Chloro-chromates,  664 
Chloro-stannates,  642 
Chromates,  662 
Chrome  alum,  66 1 
,,       green,  658 
,,       iron  ore,  657 
,,       ochre,  657 

red,  663 
,,       yellow,  663 
Chromic  anhydride,  659 
,,       chloride,  660 
,,       hydroxides,  658 
,,       sulphate,  661 
Chromite,  657 


Chromites,  661 
Chromium,  657 

anhydride,  659 

,,       chromate,  658 

,,       dioxide,  658 

,,       oxides  of,  658 

,,       sesquioxide,  658 

trioxide,  659 
Chromous  chloride,  660 

,,       hydrated  oxide,  658 

,,       sulphate,  660 
Chromyl  chloride,  663 
Chrysoberyl,  572,  618 
Cinnabar,  597 
Clark's  process  for  softening  water, 

222 

Classification  of  elements,  112 
Clay,  614 

,,       ironstone,  672 
Cleveite,  264 
Coal,  293 

,,       gas,  319 

Coarse  metal  (copper),  551 
Cobalt,  682 

,,       bloom,  682 

,,       glance,  479,  682 

,,       oxides  of,  682 
Cobaltamines,  685 
Cobaltic  hydroxide,  683 

,,       oxide,  683 
Cobalto-cobaltic  oxide,  683 
Cobaltous  chloride,  683 

,,       hydroxide,  683 

,,       oxide,  683 

,,       sulphate,  684 

,,       sulphide,  684 
Coefficient  of  absorption,  144 

,,       solubility,  144 
Coefficients  of  expansion  of  gases,  69 
Coke,  290 
Colemanite,  607 
Colloids,  636 
Columbite,  655 
Cclumbium,  655 
Combining  proportions,  30 
Combustibles,  322 
Combustion,  321 

,,       gain  in  weight  by,  325 

,,       heat  of,  326 

,,       supporters  of,  322 
Common  salt,  531 
Compound  radicals,  23 


Index 


711 


Compounds,  7 
Conductivity,  molecular,  107 
Condy's  fluid,  670 
Constant-boiling  mixtures,  155 
Constant-freezing  solution,  155 
Constant  composition,  law  of,  25,  31 
Constitution  of  matter,  3 
Contact  process,  sulphuric  acid,  432 
Copper,  550 

,,       acetylide,  318 

..       alloys,  553 

,,       arsenite,  487 

,,       bromide,  555 
carbide,  318 

,,       carbonates,  557 

,,       chlorides,  555 

,,       ferrocyanide,  156 
fluoride,  555 
glance,  550 

,,       hydride,  473 
hydroxide,  554 

,,       nitrate,  556 

,,       nitroxyl,  244 
oxides,  553 
oxychloride,  555 

,,       pyrites,  550 

,,       sulphate,  556 

,,       sulphides,  557 
Coprolites,  582 
Coral,  577 
Corpse  light,  330 
Corrosive  sublimate,  603 
Corundum,  614 
Cream  of  tartar,  497,  512 
Crith,  56 
Critical  pressure,  79 

,,       temperature,  78,  133 
Croceo-cobaltic  salts,  686 
Crocoisite,  657 
Crookesite,  624 
Cryohydric  solutions,  155 
Cryolite,  347,  614 
Crystalline  forms,  160 
Crystallisation,  suspended,  137,  151 

,,      water  of,  216 
Crystalloids,  636 
Cubic  system,  161 
Cubical  nitre,  541 
Cupel,  560 
Cupellation  process  (silver),  560 


Cupric  carbonates,  557 
,,       chloride,  555 

hydroxide,  554 
, ,       nitrate,  556 
oxide,  554 
,,      sulphate,  556 
sulphide,  557 
Cuprous  acetylide,  318 

chloride,  555 
,,       oxide,  553 
,,       sulphide,  557 
Cyanide  process  (gold),  567 

DALTON,  atomic  theory,  30 

Davy  lamp,  330 

Deacon's  process,  354  , 

Dead  Sea,  solid  matter  in,  219 

Deep  well  waters,  220 

Deliquescence,  217 

Dephlogistigated  air,  181 
,,       marine  acid  air,  352 

Dew-point,  257 

Diamidogen,  278 

Diatomic  molecules,  8 

Dialysed  iron,  678 

Dialysis,  635 

Diamond,  285 

,,       combustion  of,  287 
,,       specific  heat  of,  47 

Diffusiometer,  82 

Diffusion  of  gases,  81 

i,  >»         law  °f.  83 

,,       of    dissolved     substances, 

159 

Dimorphism,  162 
Dissociation,  88 

,,       coefficient,  108 

,,       electrolytic,  96 

,,       pressure,  94 
Disulphates,  435 
Disulphur  dichloridc,  413 
Disulphuryl  chloride,  440 
Dithionates,  437 
Divalent  elements,  59 
Dolomite,  572 
Dry  copper,  551 
Dulong  and  Petit,  law  of,  46 
Dutch  brass,  553 

,,      metal,  360 
Dyad  elements,  59 


712 


Index 


EARTH'S  crust,  composition  of,  182 

Ebullition,  129 

Efflorescence,  217 

Effusion  of  gases,  85 

Eka-aluminium,  123 

Eka-boron,  123 

Eka-silicon,  123 

Electric  furnace,  454,  615 

Electro-chemical  equivalents,  100 

Electro-gilding,  568 

Electrolysis,  96 

Electrolytes,  97 

Electrolytic  dissociation,  96,  101 

Electrons,  104 

Electroplating,  99,  563 

Elements  and  compounds,  6 

,,       classification  of,  112 

,,      list  of,  22 

,,      non-metallic.  8 
Elton  Lake,  water  or,  219 
Emerald,  572 
Emery,  617 
Empyreal  air,  181 
Endosmometer,  155 
Endosmose,  155 
Endothermic  compounds,  168 
English  brass,  553 

,,       Channel,  composition  of,  219 
Epsom  salts,  575 
Equations,  chemical,  23 
Equivalents,  chemical,  30 

„      electro-chemical,  100 
Estramadurite,  452,  582 
Ethyl  hydrogen  sulphate,  315 

,,      silicate,  631 
Ethylene,  314 

,,       dibromide,  314 
Euchlorine,  372 
Eudiometry,  252 
Evaporation,  126 

,,      cold  produced  by,  78,  130, 276 
Exothermic  compounds,  168,  331 
Expansion  by  heat  of  liquid  carbon 

dioxide,  308 
Expansion  by  heat  of  liquid  oxygen, 

193 
Extincteur,  303 

FARADAY'S  law,  100 
Felspar,  614,  637 


Ferrates,  679 
Ferric  chloride,  680 

,,       ferrocyanide,  681 

,,      hydroxide,  678 

,,  ,,        soluble,  678 

,,       oxide,  678 

,,       sulphate,  680 

,,       sulphide,  681 
Ferrites,  677 
Ferro-manganese,  674 
Ferroso-ferric  oxide,  678 

,,       sulphide,  682 
Ferrous  bromide,  520 

,,       chloride,  679 

,,       chromite,  662 

,,       ferricyanide,  680 

,,       ferrocyanide,  680 

,,       hydroxide,  677 

,,       oxide,  677 

,,       sulphate,  679 

,,       sulphide,  681 
Fettling,  675 
Fine  metal  (copper),  551 
Fire-damp,  314 
Fire-damp  caps,  331 
Fixed  air,  300 
Fixed  alkali,  506 
Flame,  332 

candle,  335 

,,       the  Bunsen,  341 

,,      structure  of,  332 
Flames     cause    of     luminosity    of, 

338 

Flint,  628 

Flintshire  furnace,  644 
Fluorapatite,  347 
Fluorides,  350 
Fluorine,  346 
Fluor-plumbates,  348 
Fluor-spar,  347,  577 
Forces,  chemical  and  physical,  3 
Formula  weight,  54 
Formulae,  23 

Fraction  of  dissociation,  the,  91 
Franklinite,  591 
Fulminating  gold,  569 

,,      silver,  564 
Fusco-cobaltic  salts,  685 
Fusible  metal,  500 
Fusion,  latent  heat  of,  138 


Index 


7I3 


GADOLINITE,  606 

Gahnite,  591 
Galena,  558,  643 
Gallium,  124,  606 
Galvanised  iron,  593 
Gas  carbon,  289 

Gases,  absorption  by  charcoal,  292 
, ,       coefficients  of  expansion  of,  69 

critical  pressure,  79 
,,       critical  temperature  of,  79 

diffusion  of,  81 
,,       effusion  of,  85 
, ,       kinetic  theory  of,  85 
,,       liquefaction  of,  72 
,,       occlusion  of,  179 
,,      relation  to  heat,  69 
,,       relation  to  pressure,  71 
,,       solubility  of,  in  liquids,  142 
,,       transpiration  of,  85 
Gastric  juice,  534 
Gay-Lussac,  law  of,  26,  38 
General  properties  of  gases,  69 

,,      liquids,  126 
German  silver,  593 
Germanium,  124,  627 
Gilding,  568 
Glauberite,  541 
Glauber's  salt,  541 
Glucinum,  572 
Gold,  567 

,,       alloys,  568 

, ,       compounds  of,  569 

fineness  of,  568 
,,       fulminating,  569 
Graduators,  532 
Graham's  law,  83 
Gramme-molecule,  57 
Graphite,  287 

,,       specific  heat  of,  47 
Greenockite,  596 
"  Green  salt  of  Magnus,"  695 
Green  vitriol,  434,  679 
Grey  antimony  ore,  498 

,,       cast  iron,  674 
Guignet's  green,  659 
Gun-cotton,  240 
Gun-metal,  553 
Gunpowder,  523 

products  of  combustion  of,  524 
Gypsum,  581 

,,       fibrous,  581 


HAEMATITE,  672 

Haemoglobin,  193 
Hair  salt,  619 
Half-electrolytes,  97 
Halogens,  18,  345 
Haloid  salts,  18 
Hardness  (water),  221 
Hargreaves'  process,  539 
Hausmannite,  666 
Heat,  atomic,  46 
,,       molecular,  49 
,,       of  combustion,  326 
,,       of  formation,  167 

specific,  45 

,,       specific,  table  of,  46 
,,      units,  165 
Heavy  spar,  586 
Helium,  267 
Henry's  law,  143 
Hepar  sulphuris,  526 
Hexagonal  system,  161 
Holmes's  signal,  464 
Horn  mercury,  602 
Horn  silver,  565 
Hydrazine,  278 

,,       hydrochloride,  279 
,,      hydrate,  279 
,,       sulphate,  278 
Hydrocarbons,  312 
Hydrofluosilicic  acid,  630 
Hydrogen,  171 

,,       atmospheric,  260 
,,       chloride,  363 

compounds      with      oxygen, 

203 

,,       dioxide,  223 
,,      disodium  phosphate,  475 
,,      displaceable,  19 
,,       liquid,  178 
,,       monoxide,  203 
,,       nitrate,  19 
,,       occlusion  of,  171,  179 
,,       peroxide,  223 
,,       persulphide,  412 
,,       phosphide,  gaseous,  460 
,,  ,,  liquid,  463 

»  »  solid,  464 

,,       position   of,  in  the  periodic 

classification,  125 
,,       potassium  fluoride,  347 
ii  ,,          sulphate,  434 

2   Y  2 


Index 


Hydrogen  sodium  ammonium  phos- 
phate, 475 
sulphate,  19 
sulphide,  408 
telluride,  448 

,,       trisulphide,  413 
Hydrogenium,  179 
Hydrolysis,  in 
Hydromagnesite,  576 
Hydroxides,  17 
Hydroxyl,  281 
Hydroxylamine,  281 

,,       disulphonate,  281 

,,       hydrochloride,  281 

,,      mono-sulphonate,  282 
Hypobismuthic  oxide,  501 
Hypobismuthous  oxide,  501 
Hypochlorites,  373 
Hypochlorous  anhydride,  371 
Hypoiodous  acid,  395 
Hyponitrous  anhydride,  248 
Hypophosphites,  472 
Hypovanadic  chloride,  656 

,,      oxide,  656 

,,       sulphate,  656 

ICE,  214 

,,       effect  of  pressure  upon,  214 

,,       the  melting-point  of,  138 
Icicle,  132 
Ignition-point,  329 
Indigo-copper,  557 
Indium,  122,  606 
Inflammable  air,  171 
International  atomic  weights,  22,  38 
Intestinal  gases,  hydrogen  in,  171 
lodates,  393 
lodic  anhydride,  391 
Iodine,  384 

,,       bromides,  396 

,,      chlorides,  396 

,,       pentoxide,  391 
Ionic  notation,  105 

theory,  the,  101 
lonisation,  107 
Ions,  99 

migration  of,  108 
Iridium,  690 

,,      chlorides,  691 

,,       oxides,  691 

Irish  Sea,  solid  impurity  in,  219 
Iron,  671 


Iron  alum,  68 1 

,,       carbide,  513,  674 
carbonyl,  300 

,,       magnetic  oxide  of,  678 

,,       monoxide,  677 

,,       oxides  of,  677 

,,       passive,  677 

,,       pyrites,  400,  672 

,,       sesquioxide,  678 

,,       sesquisulphide,  681 

,,       sulphides  of,  681 
Isodimorphism,  162 
Isogonism,  52 
Isomerism,  194 
Isometric  system,  161 
Isomorphism,  51,  162 
law  of,  51 

JOLLY'S  apparatus,  255 

KAINITE,  512 
Kelp,  385 

Kelp  substitute,  386 
Kiesel-guhr,  633 
Kieserite,  572,  575 
Kinetic  theory,  85 
Kish,  288 
Kryptcn,  269 
Kupfernickel,  479,  687 

LAGOONS  (boric  acid),  609 

Lakes,  618 

Laminaria  digitata,  385 
,,       stenophylla,  385 

Lamp-black,  289 

Lanarkite,  643 

Lanthanum,  606 

Latent  heat  effusion,  138 

,,         ,,         vaporisation,  130 

Laughing-gas,  248 

Law  of  Boyle,  71 
Charles,  69 
constant  heat  consummatidh, 

169 

„       constant  proportion,  25,  26 
,,       Dulong  and  Petit,  45 
,,       gaseous  diffusion,  83 
,,       Gay-Lussac,  26,  38 
,,       multiple    proportions,     25, 

27 
,,       octaves,  113 


Index 


715 


Law  of  partial  pressures,  147 

,,       periodic,  112 

,,       reciprocal  proportions,  25,  28 
Layer  crystals,  52 
Lead,  643 

,,       acetate,  652 

,,       action  of  water  upon,  646 

,,       carbonate,  651 

,,       chromate,  663 

,,       compositionof  commercial, 647 

,,       desilverisation  of,  561 

,,       dichloride,  649 

,,       dioxide,  649 

,,       disulphate,  653 

,,       ethide,  628 

,,       nitrate,  651 

,,       oxides  of,  647 

,,       oxychloride,  650 

,,       sesquioxide,  648 

,,       softening  of,  645 

,,       squirted,  647 

,,       suboxide,  647 

,,       sulphate,  653 

,,       sulphide,  654 

,,       sulphochlorides,  654 

,,       tetracetate,  654 

,,       tetrachloride,  650 

,,      tree,  645 

,,       white,  651 
Leblanc  process,  534 
Leguminous  plants,  258 
Lepidolite,  543 
Light  red  silver  ore,  558 
Lime,  578 

chloride  of,  580 

,,       dead  burnt,  578 

, ,       milk  of,  579 

,,       quick,  578 

,,       slaked,  578 

, ,       superphosphate  of,  583 
Limestone,  577 
Lines  of  symmetry,  160 
Liquefaction  of  air,  77 

,,      of  gases,  72 

Liquids,  general  properties  of,  126 
Liquor  ammonics,  275 
Litharge,  647 
Lithium,  543 

,,      carbonate,  544 

,,       hydroxide,  544 

,,       mica,  543 


Lithium  nitride,  232 

,,       oxide,  544 

,,       phosphate,  544 

,,       spectrum  of,  509 
Liver  of  sulphur,  526 
Load-stone,  672,  678 
Lothar  Meyer's  curve,  120 
Lucifer  matches,  459 
Luminous  paint,  584 
Lunar  caustic,  566 
Luteo-cobaltic  salts,  686 

MAGISTRAL,  559 
Magnesia,  574 
Magnesia  alba  levis,  576 

,,      ponder osa,  577 

,,       usta,  574 
Magnesia  mixture,  575 
Magnesian  limestone,  220,  577 
Magnesite,  576 
Magnesium,  572 

,,      aluminate,  618 

,,       ammonium  chloride,  575 

,,       ammonium  phosphate,  475 

,,       boride,  613 

,,      bromate,  384 

,,       calcium  chloride,  575 

,,       carbonates,  576 

,,       chloride,  574 

,,       combustion  of,  in  steam,  174 

,,       hydroxide,  574 
nitride,  232,  574 

,,      oxide,  574 

,,       oxychloride,  575 

,,       phosphate,  475 

,,       platinocyanide,  217 

, ,       potassium  chloride,  57 

,,       pyrophosphate,  476 

,,       silicide,  629 

sulphate,  575 
Magnetic  iron  ore,  672 

,,       oxide  of  iron,  678 

,,       pyrites,  682 
Magnetite,  678 
Malachite,  550 
Manganates,  669 
Manganese,  666 

blende,  666 

,,       dioxide,  667 

,,       monoxide,  666 

,,      oxides  of,  666 


716  Index 

Manganese  sesquioxide,  667 

,,       spar,  666 
Manganic  chloride,  668 

,,       oxide,  667 

,,      sulphate,  668 
Manganite,  666 
Manganites,  668 
Mangano-manganic  oxide,  667 
Manganous  chloride,  668 

,,      chroraite,  662 

,,       hydroxide,  667 

,,       sulphate,  668 
Marble,  582 
Marine  acid  air,  363 
Marsh  gas,  312 

,,       synthesis  of,  443 
Marsh's  test,  482 
Massicot,  647 
Malches,  459 
Matlockite,  643 
Mechanical  mixtures,  8 
Mediterranean  Sea,  219 
Meerschaum,  573 
Mendelejeff  s  periodic  law,  113 
Mephitic  air,  229 
Mercuric  ammonium  chloride,  605 

,,      chloride,  603 
.     ,,      iodide,  603 

,,      oxide,  601 

,,       potassium  chloride,  67 
Mercurius  calcinatus  per  se,  181 
Mercurous  chloride,  602 
nitrate,  602 

,,       oxide,  601 

,,      sulphate,  602 
Mercury,  597 

,,      alloys  of  (amalgams),  600 

, ,       deadening  of,  600 

,,       distillation  of,  598 

,,      oxides  of,  601 
Metal  slag  (copper),  551 
Metallic  carbonyls,  299 

,,      nitroxyls,  244 
Metalloids,  8 
Metals  and  non-metals,  7 
Metameric  compounds,  194 
Metantimonates,  497 
Metaphosphates,  476,  498 
Metarsenates,  488,  498 
Metarsenites,  487 


Metastannates,  641 

Metavanadates,  655 

Meteoric  iron,  171 

Methane,  312 

Meyer,     Lothar,     curve    of    atomic 

volumes,  120 
Microcosmic  salt,  543 
Migration  of  ions,  108 
Milk  of  lime,  579 

,,      sulphur,  407 
Milky  quartz,  634 
Mineral  alkali,  506 
Minium,  648 
Mispickel,  479 
Mixed  crystals,  52 
Modes  of  chemical  action,  13 
Molecular  combinations,  67 
conductivity,  107 

,,       concentration,  94 

, ,       depression  of   the   freezing  - 
point,  140 

, ,       elevation  of  the  boiling-point, 

134 
,,       equations,  55 

formulae,  23 
,,       heats,  49 
,,       lowering  of  vapour  pressure, 

,,       volume,  44 

,,       weight,  41 

,,       weight,  determination  of,  by 
the  depression  of  freezing, 
point,  140 
Molecules,  3 

,,       compound,  6 

,,       definition  of,  4 

,,       elementary,  6 

,,       mean  free  path  of,  86 

,,       size  of,  3 
Molybdates,  664 
Molybdenite,  664 
Molybdenum,  664 

,,      chlorides,  665 

,,      ochre,  664 

,,       oxides,  664 
Monad  elements,  59 
Mono-atomic  molecules,  8 
Monoclinic  system,  162 
Monosymmetric  system,  162 
Monovalent  elements,  59 


Index 


717 


Mordants,  618 
Mortar,  579 

,,       the  setting  of,  579 
Mosaic  gold,  643 
Mottramite,  655 
Mundic,  479 
Muntz  metal,  553 
Multiple  proportions,  law  of,  25,  31 

NATURAL  waters,  218 
Natural  steel,  624 
Neon,  269 

Nessler's  solution,  272,  605 
Neutral  alum,  622 
Nickel,  687 

alloys  of,  688 
blende,  687 

,,  carbonyl,  299 
chloride,  689 
glance,  479,  687 

,,       monosulphide,  689 

,,      monoxide,  688 
oxides  of,  688 

,,       sesquioxide,  688 

,,       silver,  593 

sulphate,  689 

Nickelo-nickelic  oxide,  689 
Nickelous  oxide,  688 

,,      sulphide,  689 
Niobates,  655 
Niobium,  655 

,,       oxides  of,  655 
Nitrates,  241 

,,       detection  of,  241 
Nitre,  522 

plantations,  523 

Nitric  acid,  manufacture  from  atmos- 
pheric nitrogen,  235 

,,       anhydride,  241 

,,       oxide,  246 
Nitrides,  278  280 
Nitrification,  522 
Nitrites,  245 
Nitro-cellulose,  240 
Nitrogen,  229 

,,       iodide,  283 

,,       oxides  and  oxyacids  of,  234 

,,       oxyfluorides,  251 

,,       pentoxide,  241 

,,       peroxide,  242 

,,       tribromide,  283 

,,       trichloride,  282 
Nitro-metals,  244 


Nitro-sulphuric  acid,  426 
Nitrosyl  chloride,  250 
fluoride,  251 

,,       hydrogen  sulphate,  251 

,,       sulphate,  426 
Nitroxyl  fluoride,  251 
Nitrous  anhydride,  234 
Noble  metals,  240 
Nomenclature,  15 

,,       of  ions,  105 
Non-electrolytes,  97 
Non-metals,  7 
"  Nordhausen  "  acid,  434 
Notation,  chemical,  21,  53 


OCCLUDED  hydrogen,  179 

Occlusion  of  gases,  171 

Olefiant  gas,  314 

Opal,  633 

Ore  hearth,  644 

Orangeite,  627 

Organic  chemistry,  definition  of,  296 

Orpiment,  479 

Orthite,  606 

Orthoclase,  637 

Orthorhombic  system,  161 

Osmiridium,  690 

Osmium,  690 

,,       oxides  of,  691 

,,       tetroxide,  691 
Osmotic  pressure,  155 
Osteolite,  582 
Oxides,  17 
Oxygen,  181 

,,       allotropic,  195 

,,       Brin's  process,  184 

,,       Tessie"  du  Motay  process,  189 
Oxyhaemoglobin,  193 
Oxyhydrogen  flame,  327 
Oxymuriatic  acid,  352 
Ozone,  195 

,,       atmospheric,  256 

, ,       constitution  of,  199 

,,       tube,  Siemens',  195 

,,  ,,     Andrews',  200 

PALLADIUM,  690 

,,       absorption  of  hydrogen  by, 

179 

, ,       chlorides,  691 
,,       hydride,  179 
,,       oxides,  691 


718 


Index 


Parkes's  process,  561 
Partial  pressures,  law  of,  147 
Partially  miscible  liquids,  149 
Passive  iron,  677 
Pattinson's  process,  561 
white  lead,  650 
Pearl-ash,  521 
Per  chlorates,  375 
Percy-Patera  process,  562 
Perdisulphuric  acid,  424 
Periclase,  574 
Peridote,  637 
Periodates,  393 
Periodic  classification,  112 
Permanent  white,  589 

,,       hardness,  221 
Permanganates,  669 
Permanganic  anhydride,  670 
Permonosulphuric  acid,  425 
Persulphates,  425 
Persulphuric  anhydride,  424 
Petalite,  543 
Petzite,  567 
Pewter,  639 
Phenacite,  572 
Phlogiston,  321 
Phosgene  gas,  299 
Phospham,  477 
Phosphates,  474 
Phosphine,  460 
Phosphites,  473 
Phosphonium  bromide,  462 

,,       chloride,  462 

,,       iodide,  462 
Phosphoretted   hydrogen,    gaseous, 

460 

Phosphorous  oxide,  470 
Phosphorus,  451 

,,       allotropic,  458 

,,       compounds     with      sulphur, 
478 

,,       manufacture  of,  452 

,,  ,,         by  electric  furnace, 

454 

,,       oxides  and  oxyacids,  469 
,,       oxychloride,  468 
,,       oxyfluoride,  468 
,,       pentabromide,  467 
,,       pentachloride,  465 
pentafluoride,  464 
,,       pentasulphide,  478 


Phosphorus  pentoxide,  471 

red,  458 

,,       tetriodide,  467 
,,       tribromide,  467 
,,       trichloride,  465 
,,       trifluoride,  464 

triodide,  467 
Phosphoryl  chloride,  468 

fluoride,  468 
,,       nitride,  478 

triamide,  478 
Photo-salts,  566 
Physical  constants  of  gases,  80 
Pig-boiling,  675 
Pig  iron,  674 
Pitchblende,  664 
Planes  of  symmetry,  160 
Plaster  of  Paris,  581 
Plastic  sulphur,  406 
Plate  sulphate,  386 
Platinamines,  695 
Platinates,  693 
Platinic  hydroxide,  693 

,,       chloride,  694 
Platiniridium,  690 
Platino-chlorides,  693 
,,       cyanides,  695 

nitrites,  695 

Platinotype  process,  694 
Platinous  chloride,  693 
,,       hydroxide,  693 
Platinum,  691 
,,       alloys,  693 
,,       black,  693 

oxides  of,  693 
oxysalts,  695 
,,       sodium  chloride,  67 
,,       spongy,  692 
,,       sulphides  of,  695 
,,       tetrachloride,  694 
Platoso-ammonium  compounds,  695 
Plumbago,  288 
Plumbic  chloride,  649 
,,       oxalate,  647 
,,       oxide,  647 
,,       peroxide,  649 
Plumbous  oxide,  647 
Plumbum  nigrum,  643 
Pollux,  510 
Polybasite,  558 
Polyhalite,  520 


Index 


719 


Polymerism,  194 
Pot-ashes,  521 
Potash,  caustic,  515 
Potassium,  510 

,,       alum,  621 

,,       aluminate,  618 

,,       antimonate,  498 

,,       borofluoride,  608,  612 

,,       bromate,  384,  520 

,,       bromide,  520 

,,       carbonate,  521 

,,       chlorate,  516 

,,  ,,         electrolytic      manu- 

facture of,  518 

,,       chloride,  516 

,,       chlorochromate,  664 

,,       chloroplatinate,  694 

,,       chloroplatinite,  694 

,,       chromate,  662 

,,       chromium  alum,  620,  661 

,,       dichromate,  662 

,,       ferrate,  679 

,,       ferricyanide,  68 1 

,,       ferrocyanide,  297,  680 
fluoride,  515 

,,       fluor-plumbate,  348 

,,       hydride,  514 

,,       hydroxide,  515 
hypoiodite,  283 

,,       hyponitrite,  250 
iodate,  393 

,,      iodide,  520 

manganate,  669 
metaborate,  610 

,,       metantimonate,  498 
metarsenite,  487 

,,       metastannate,  641 

,,       meta-thio-arsenitc,  490 

,,       nitrate,  522 

,,       nitrite,  245 

,,       osmate,  691 

,,       oxides  of,  514 

,,      ortho-thio-antimonate,  499 

,,       ortho-thio-antimonite,  499 

,,       ortho-thio-arsenate,  490 

,,       ortho-thio-arsenite,  490 
pentasulphide,  525 

.,       pentathionate,  439 

,,       perchlorate,  519 

.,      periodate,  394 


Potassium  permanganate,  670 
,,       peroxide,  514 

platinic  chloride,  694 
,,       platino-cyanide,  695 
, ,       platinous  chloride,  694 
„       plumbate,  649 
,,       pyro-antimonate,  498 
,,      ruthenate,  691 
,,       silico-fluoride,  629 
,,      silver  thiosulphate,  437 
,,       stannate,  640 
,,       sulphate,  520 
,,       sulphite,  421 
,,      sulphides  of,  525 
,,      tetrachromate,  663 
-,,      trichromate,  663 
,,       zinc  oxide,  175 

Powder  of  Algaroth,  495 

Praseo-cobaltic  salts,  685 

Preparing  salt,  640 

Producer  gas,  186 

Proustite,  558 

Prussian  blue,  681 

Pseudo-alums,  620 

Pucherite,  655 

Puddling,  675 

Purple  copper  ore,  550 

Purpureo-cobaltic  salts,  686 

Pyrargyrite,  558 

Pyrites  burners,  428 

Pyrolusite,  666 

Pyromorphite,  643 

Pyrophosphates,  476 

Pyrosulphuric  chloride,  440 

QUANTITATIVE  notation,  53 
Quartz,  633 
Quicklime,  578 

RADIATED  pyrites,  681 

Radicals,  compound,  23 

Radium,  697 

Rain  water,  solid  impurity  in,  220 

Raoult's  method,  140 

Realgar,  489 

Red  antimony,  491 

,,      copper  ore,  553 

,,       haematite,  672 

,,      lead,  648 

,,      manganese  oxide,  667 


72O 


Index 


Red  phosphorus,  458 
,,       zinc  ore,  591 
Refinery  slag  (copper),  551 
Regular  system,  161 
Reiset's  second  base,  chloride  of,  695 
Relation  of  gases  to  heat,  69 

»>  M         i»     pressure,  71 

Reversible  reactions,  88 
Rhodium,  690 
Rhombic  system,  161 
Rochelle  salt,  563 
Rock  crystal,  633 
Rock  salt,  526 
Rodonda  phosphates,  452 
Roll  sulphur,  403 
Roman  alum,  621 
Roseo-cobaltic  salts,  686 
Rouge,  678 
Rubidium,  544 
Rubies,  artificial,  618 
Ruby,  617 
Ruby  ore,  550 

,,     silver  ore,  558 

,,     sulphur,  479 
Rust  of  iron,  676,  678 
Ruthenium,  691 

,,       chlorides  of,  691 
,,       oxides,  691 
Rutile,  627 

SAL  alembroth,  603 

,,       ammonia,  546 
Salt-cake  process,  534 
Salt-forming  oxides,  17 
Salterns,  531 
Saltpetre,  522 
Salts,  acid,  19 

,,       basic,  20 

,,       haloid,  18 

,,       normal,  19 

,,       of  hydrogen,  107 

,,       ofhydroxyl,  107 

,,       oxy-,  18 

,,       thio-,  18 
Sand,  628 
Sapphire,  617 
Satinspar,  581 
Saturated  solutions,  151 

,,       vapours,  127 
Scandium,  606 


Scheele's  green,  487 
Scheelinite,  664 
Schlippe's  salt,  499 
Schonite,  576 
Schweinfurt  green,  487 
Scotch  hearth,  644 
Seaweed,  iodine  in,  385 
Selenite,  581 
Selenium,  444 

,,       alums,  620 

,,       dichloride,  446 

,,       dioxide,  447 
Selenuretted  hydrogen,  446 
Seltzer  water,  220 
Semipermeable  membranes,  156 
Serpentine,  573,  637 
Siemens'  ozone  tube,  195 
Silica,  633 
Silicates,  636 

Siliciuretted  hydrogen,  630 
Silicon,  628 

,,       chloride,  632 

,,       chloroform,  628 

,,       dioxide,  633 

,,       fluoride,  632 

,,       hexachloride,  633 

,,       hexafluoride,  632 

,,       hydride,  630 

,,       liquid,  631 
Silver,  558 

,,       allotropic,  563 

,,       alloys,  563 

,,       alum,  566 

,,       bromide,  565 

,,       chloride,  564 

,,       flashing  of,  560 
fluoride,  565 

,,       fulminating,  564 
glance,  558 
iodide,  565 

,,       nitrate,  566 
oxides,  563 

,,       oxybromide,  566 

,,       oxychloride,  566 

,,      periodate,  394 

phosphates,  475,  477 
plating,  563 
spitting  of,  563 

,,       standards,  563 

,,       suboxide,  564 


Index 


721 


Silver  sulphate,  522,  566 

,,       sulphide,  558 
Slaked  lime,  578 
Smalt,  686 
Smaltine,  682 
Smoky  quartz,  634 
Soda,  540 
Soda-ash,  538 

caustic,  530 
,,       crystals,  540 
Sodium,  526 

acetate,  313 

,,       alloy  with  potassium,  529 

,,       aluminate,  614 

,,       aluminium  chloride,  615 

,,       amalgam,  601 

,,       antimonate,  497 

,,       antimonite,  496 

,,       arsenate,  488 

,,      benzoate,  279 

,,       bicarbonate,  540 

„      bromide,  534 

„      carbonate,  534 

,i  ,,        electrolytic    manu- 

facture of,  539 

,,       chloride,  531 

,,       chloro-platinate,  694 

„      electrolytic    manufacture  of, 
526 

j.       electrolytic     manufacture    of 
(Borchers*  process),  527 

t,      hydrazoate,  279 

,,       hydride,  529 

,,       hydroxide,  530 

.,       hypophosphite,  473 

,,       hyposulphite,  436 

,,       iodide,  534 

,,      metabisulphite,  424 

,,      metaniobate,  655 

,,       metaphosphate,  476 

,,       metastannate,  641 

,,       metatantalate,  655 

,,       metavanadate,  655 

,,       nitrate,  541 

,,       nitride,  280 

,,       oxalate,  175 
oxides,  529 

,,       permanganate,  670 

,,      phosphates,  542 


Sodium  pyro-arsenate,  488 

,,       pyrophosphate,  476 

,,       sesquicarbonate,  540 
silicate,  635 

,,       silver  thiosulphate,  562 

, ,       stannite,  640 

,,      sulphate,  541 

,,  ,,        solubility  curve,  153 

,,       sulphide,  534 

,,       thio-antimonate,  499 

,,       thiosulphate,  436 
tungstate,  664 

,,       uranate,  664 

zinc  chloride,  66 
Soffioni,  609 
Solar  prominences,  171 
Solder,  639 
Solfatara,  399 

Solidification,  suspended,  137, 404, 456 
Solidifying  points  of  liquids,  137 

,,       points  of  liquids,  effect  of  dis- 
solved substances  upon,  139 

,,       points    of   liquids,    effect    of 

pressure  on,  137 
Solubilities,  diagram  of,  152 
Solubility  of  gases  in  liquids,  142 

,,       of  liquids  in  liquids,  148 

,,       of  mixed  gases,  146 

,,       of  solids  in  liquids,  150 
Solution,  142 
Solutions,  saturated,  151 

,,       supersaturated,  151 
Sombrerite,  452,  582 
Spathic  iron  ore,  672 
Specific  gravity  of  gases,  40 

,,  ,,       liquids  and  solids,  119 

„       heat,  45 

,,      heats,  tables,  46 
Spectra  of  alkali  metals,  505 
Spectroscope,  507 
Specular  iron  ore,  672 
Speiss-cobalt,  682 
Spiegel,  674 
Spinelle,  618 
Spirits  of  hartshorn,  272 
Spitting  of  silver,  the,  565 
Spodumene,  543 
Spring  water,  219,  220 
Stalactites,  222 
Stalagmites,  222 


722 


Index 


Standard  temperature  and  pressure, 

69,  71 

Stan  nates,  640 
Stannic  chloride,  642 

,,       sulphide,  642 
Stannous  chloride,  641 

hydrated  oxide,  639 

,,       nitrate,  638 

,,       oxide,  639 

oxychloride,  641 

,,       sulphate,  638 

,,       sulphide,  642 
Stassfurt  deposits,  512,  520,  574 
Steam,  214 

,,      volume,  composition  of,  208 
Steel,  675 
Steel  mill,  329 
Stephanite,  558 
Stereotype  metal,  492 
Stibnite,  498 

Still-liquor,  composition  of,  357 
Stream-tin,  637 
Stromeyerite,  558 
Strontia,  584 
Strontianite,  584 
Strontium,  584 

ammonium,  584 

,,  chloride,  585 
dioxide,  585 
hydride,  584 

,,       hydroxide,  584 

,,       nitrate,  585 

,,       oxides,  584 

,,       sulphate,  585 
Substitution,  382 
Suint,  511 
Sulphates,  434 
Sulphides,  410 
Sulphion,  105 
Sulphites,  421 
Sulpho-acids,  17 
Sulpho-thionyl  chloride,  414 
Sulphovinic  acid,  315 
Sulphur,  398 

,,       allotropic  modifications,  404 

,,       chlorides  of,  413 

,,  dioxide,  415 
flowers  of ,  402 

,,       milk  of,  407 

,,       oxides  and  oxyacicls  of,  414 

,,       oxychlorides  of,  439 

,,       perfluoride,  441 


Sulphur,  plastic,  406 

,,       prismatic,  404 

,,       recovery  of,  from  alkali- waste, 
400 

, ,       recovery  of  (Chance's  process), 
411 

,,       rhombic,  404 

,,       sesquioxide,  423 
tetrachloride,  '414 

,,       trioxide,  421 
Sulphuretted  hydrogen,  408 
Sulphuric  acid,  contact  process,  432 
, ,          , ,      manufacture  of,  428 
Sulphuric  anhydride,  414 
,,     chlorhydrate,  440 
Sulphurous  anhydride,  414 
Sulphuryl  chloride,  439 
Supercooling  of  water,  136 
Superphosphate  of  lime,  583 
Supersaturated  solutions,  151 
Suspended  solidification,  137,  404,  456 
Sylvanite,  567 
Sylvine,  512 
Symbols,  21 
Sympathetic  inks,  217 
Synthesis,  13 

TACHYDRITE,  575 
Talc,  573 
Tank  liquor,  537 
Tantalite,  655 
Tantalum,  655 

,,       oxides  of,  655 
Tartar  emetic,  497 
Tellurates,  449 
Telluretted  hydrogen,  448 
Tellurites,  449 
Tellurium,  448 
Temporary  hardness,  221 
Tenorite,  554 

Tessie"  du  Motay  process,  189 
Tetradymite,  448 
Tetragonal  system,  161 
Tetratomic  molecules,  8 
Tetravalent  elements,  59 
Thallic  chloride,  625 

,,       nitrate,  626 
oxide,  607,  625 
sulphate,  626 

,,       sulphide,  607 
Thallium,  623 

,,       oxides  of,  624 


Index 


723 


Thallium  oxyhydroxide,  625 

,,       perchlorate,  607 

,,       sulphate,  607 
Thallous  carbonate,  626 

,,       chloride,  625 

,,       hydroxide,  624 

,,       iodide,  607 

,,       oxide,  624 

,,      phosphate,  626 
Thenardite,  541 
Thermochemistry,  163 
Thio-acids,  17 
Thio-antimonates,  499 
Thio-antimonites,  499 
Thio-arsenates,  490 
Thio-arsenites,  490 
Thiocarbonates,  443 
Thionyl  chloride,  439 
Thiophosphoryl  chloride,  469 

,,       fluoride,  468 
Thorite,  627 
Thorium,  627 
Tin,  637 

,,       alloys  of,  639 

,,       dioxide,  640 

,,       oxides  of,  639 

,,       oxymuriate,  642 
Tin-plate,  639 
Tin-stone,  637 
Tin-white  cobalt,  479 
Tincal,  607 
Tinning,  639 
Titanium,  627 
Tombac,  553 

Transitional  elements,  115,  671 
Transpiration  of  gases,  85 
Triad  elements,  59 
Triclinic  system,  162 
Tridymite,  633 
Triethylamine,  150 
Triethyl  silico-formate,  631 
Trivaleut  elements,  59 
Trona, 540 

Truncated  crystals,  162 
Tungstates,  664 
Tungsten,  664 

,,       chlorides,  665 

,,       oxides,  664 
Turnbull's  blue,  680 
Turpeth  mineral,  434 


Turquoise,  614 
Type  metal,  492 
Typical  elements,  115 
Twin  crystals,  634 

ULEXITE,  607 

Unit  of  heat,  165,  326 

,,       volume,  44 
Unsaturated  compounds,  62 
Uranates,  664 
Uraninite,  267 
Uranium,  664 

,,       chlorides,  665 

oxides,  664 
Uranous  salts,  665 

,,      sulphate,  665 
Uranyl  salts,  665 
Urea,  13,  24,  295 

VALENCY,  59 
Vanadates,  655 
Vanadite,  655 
Vanadium,  655 

,,       chlorides  of,  656 

,,      oxides  of,  655 

,,       oxy chlorides  of,  656 
Vaporisation,  latent  heat  of,  130 
Vapour  densities  of  elements,  42 

,,       pressures  of  solutions,  133 
Vapour  tension,  128 
Verdigris,  557 
Vermilion,  604 
Vinasse,  521 

,,       cinder,  521 " 
Vital  force,  295 
Vitriol  chambers,  430 
Volatile  alkali,  506 

WATER,  203 

,,       Clark's  process  for  softening, 

222 
,,       colour  of,  212 

compressibility  of,  213 
, ,       electrolysis  of,  207 
,,       freezing  of,  131 
,,       gas,  297 
,,       gravimetric    composition    of, 

210 
,,       hardness  of,  221 

maximum  density  of,  214 


724 


Index 


Water  of  constitution,  218 

of  crystallisation,  216 
,,       rain,  220 

solubility  of  gases  in,  147 

,,        ,,  salts  in,  150 
solvent  power  of,  216 
,,       supercooling  of,  136 

volumetric     composition     of, 

206 

Waters,  chalybeate,  220 
,,       dangerous,  223 
,,       deep  well,  220 
,,       fresh,  220 
,,       hard,  221 
,,       mineral,  219 
natural,  218 
,,       potable,  222 
river,  220 
safe,  223 
,,       sea,  219 
,,       spring,  219 
,,       suspicious,  223 
Wavellite,  452 
Weldon's  process,  357 
Welsbach  burner,  339 
White  arsenic,  485 
,,       cast  iron,  674 
,,       lead,  651 

metal  (copper),  551 
„       nickel,  687 
,,       vitriol,  218 
Witherite,  586 
Wohlerite,  627 
Wolfram,  664 
,,      ochre,  664 


Wood's  fusible  metal,  500 
Wrought  iron,  675 
Wulfenite,  664 
Wurtzite,  595 

XANTHO-COBALTIC  salts,  686 
Xenon,  269 

YTTERBITE,  606 
Ytterbium,  606 
Yttrium,  606 

ZIERVOGEL  process,  561 
Zinc,  591 

, ,       alloys  of,  593 
,,       aluminate,  591 
,,       amalgam,  601 
,,       blende,  591 
carbonate,  596 
chloride,  594 
,,       chromite,  662 
,,       granulated,  174 
,,       hydroxide,  594 

methyl,  313 
,,       nitrate,  240 
oxide,  593 
spar,  591 

,,       spinnelle,  591 
sulphate,  595 
sulphide,  595 
,,       white,  594 
Zsinci  carbonas,  596 
Zinc-copper  couple,  173,  313 
Zircon,  627 
Zirconium,  627 


Printed  by  BALLANTYNE,  HANSON  &*  Co. 
Edinburgh  6f  London 


